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Inorganic  Chemistry 
Syllabus 

BY 

HUBERT  C.  CAREL,  B.  S., 

Instructor  in  Medical  Chemistry 
UNIVERSITY  OF  MINNESOTA. 


PUBLISHED  BY 

University  Book  Store 

MINNEAPOLIS 

1897 


rVH  * • 


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I 8 8 3 X 


C\?A 


INTRODUCTORY  DEFINITIONS. 

Science — 

Classified  knowledge  and  deduced  relation;  natural  science 
objective  not  subjective. 


Divided — 

Into  two  classes  which  shade  into  each  other: 

Chemistry 

Inanimate 

Zoology 


Animate  j Biology 


Botany 


Physics 

Mineralogy,  etc. 


Physics  and  Chemistry — 

Treat  of  changes  of  matter;  i.  e.,  that  which  occupies  space. 

Physical — 

Change  in  place,  condition  or  properties,  substance  un- 
changed. 


Chemical — 

Change  of  substance  and  properties. 

Chemistry — 

Treats  of  what  substances  are  composed,  what  occurs  when 
they  change  and  the  dependence  of  properties  on  composition- 
chemical  knowledge  depends  on  analysis,  and  synthesis. 

Analysis — 

Chemical  decomposition  of  complex,  to  simple  substances. 

Synthesis — 

Chemical  combination  of  simple  to  complex  substances. 

Element — 

A substance  as  yet  undecomposed.  All  matter  is  formed 
from  one  element  or  a combination  of  elements. 


Allotropic — 

Modification  of  an  element  is  a substance  of  same  elementary 
composition  but  differing  in  chemical  and  physical  properties. 

Isomerism  and  Isomers — 

The  above  phenomenon  in  compounds  is  termed  isomerism 
and  such  bodies  are  isomers. 


Physical — 

Form  of  elements— gases,  liquids  and  solids. 

Chemical — 

Division — metals  and  non-metals. 

Metal  (in  general) — 

Chemical  element  which  unites  with  the  elements  of  water 
to  form  a base  (p.  15).  Except  hydrogen,  metals  in  general  are 
opaque  solids,  more  or  less  malleable,  ductile  and  tenacious, 
conductors  of  heat  and  electricity  and  possessed  of  a peculiar  lus- 
tre termed  metallic. 

Non=Metals  (in  general)  — 

Are,  as  the  name  suggests,  the  direct  opposites  of  metals, 
i.  e.,  they  unite  with  elements  of  water  to  form  an  acid.  Three 
of  the  non-metals  are  gases  : Oxygen,  Nitrogen,  and  Chlorine — 
one  probably  a gas,  Fluorine — oneisa  liquid,  Bromine — the  rest 
are  solids. 

Compound- 

Union  of  two  or  more  elements  in  simple  proportions,  to 
form  a new  substance. 

Mixture — 

Mechanical  intermingling  of  matter  in  any  or  all  propor- 
tions, without  change  of  substance.  The  original  materials 
may  be  recovered  by  mechanical  means. 

Chemical  Affinity — 

Tendency  of  elements  to  unite — C.  A.  of  a given  element  may 
vary  in  uniting  with  different  elements,  but  is  always  constant 
with  any  particular  element.  Some  elements  unite  directly  to 
form  a compound  A+B=AB.  More  often  an  element  unites 
with  a compound  A+CD=AD+C;  or  two  compounds  react  to 
form  two  new  compounds  AB-j-CD=AC-l-BD. 

Metathesis— 

The  last  reaction  is  known  as  metathesis  (Gr.,  to  set  over), 
and  occurs  especially  in  solution,  where  one  of  the  produces  is 
insoluble,  or  by  heat,  where  one  of  the  products  is  volatile. 

Heat  and  Chemical  Energy — 

Save  nitrogen,  all  union  of  elements  yields  heat;  that  is, 


chemical  energy  transforms  to  heat.  Decomposition  of  com- 
pounds absorbs  heat  which  is  transformed  into  chemical  energy. 

Acid  (in  general) — 

Compound  of  hydrogen  which  yields  salts  by  replacing  its 
hydrogen  with  a metal.  Acid  reddens  blue  litmus  and  possesses 
characteristic  “acid”  taste.  That  part  of  acid  which  unites 
with  the  metal  is  called  “acid  radical.” 

Base  (in  general) — 

Oxide  or  hydroxide  (OH)  of  a metal  which  exchanges  its 
metal  for  the  hydrogen  of  an  acid  forming  a salt.  Base  blues 
litmus  and  possesses  characteristic  “alkali”  taste. 

Salt— 

The  generally  neutral  result  of  the  combination  of  metal  of 
base  with  radical  of  acid. 

(base)  (acid)  (salt) 

2NaOH  + H2S04  = Na2S04  + 2H20 

(met-  (acid  (met-  (acid 

al)  radical)  al)  radical) 

Neutralization — 

This  reaction  of  an  acid  body  with  alkali  forming  a sub- 
stance neither  acid  nor  alkaline  is  termed  neutralization. 

LAWS  OF  CHEMICAL  PROPORTIONS. 

Composition  of  a given  substance  is  invariable. 

Definite  Proportions — 

Chemical  combination  always  takes  place  between  definite 
masses.  Any  excess  is  unacted  on. 

Multiple  Proportions — 

When  elements  unite  in  more  than  one  proportion,  the  ratio 
of  succeeding  compounds  are  simple  multiples  of  the  first  ratio. 
Example : N20 — N202 — N2C>3— N204 — N20s. 

Atomic  Theory. 

Above  laws  depend  on  the  Atomic  Theory  that  every  body 
is  an  aggregate  of  atoms  or  ultimate  indivisible  particles.  The 
absolute  weight  of  an  atom  is  of  course  unobtainable,  but  we 
have  relative  atomic  weights  with  hydrogen  as  the  unit. 


Atom  vs.  Molecule. 

Atom  = smallest  indivisible  particle. 

Molecule  = smallest  divisible  particle. 

Law  of  Charles — 

Volume  of  gases  varies  directly  as  the  absolute  temperature. 

Law  of  Mariotte — 

Volume  of  gases  varies  inversely  as  pressure. 

Law  of  Avagadro — 

At  same  temperature  and  pressure  equal  volumes  of  all  gases 
contain  the  same  number  of  molecules. 

Certain  Physical  Properties  — 

Of  bodies  are  always  the  same  for  the  same  substance,  i.  e.: 

(1)  The  boiling  point  of  liquids. 

(2)  The  melting  point  of  solids. 

(3)  The  crystalization  of  solids. 

SOLUTION. 

Some  substances  require  particular  solvents,  as  ether  for 
fats,  carbon  disulphide  for  yellow  phosphorus,  etc.,  but  the 
great  solvent  is  water.  In  stud3ung  the  phenomena  of  solution 
we  notice,  (1)  the  uniform  distribution  of  dissolved  substance, 
(2)  the  almost  infinite  subdivision  of  material,  (3)  that  sub- 
stances vary  as  to  solubility,  and  that  the  presence  of  foreign 
matter  often  affects  solution. 

(4)  Evaporate  a solution  and  (a)  solidsremain,  (b)  gases 
escape,  (c)  liquids  pass  off  at  their  boiling  points. 

(5)  Solution  is  of  great  advantage  in  chemical  reactions 
where  often  the  new  substance  is  insoluble  and  can  be  separ- 
ated. 

GROUPING  OF  ELEMENTS  WITH  REGARD  TO 
PROPERTIES. 

Non=Meta!s — 

Group  I— Cl,  Br,  I,  F. 

Group  II — 0,  S,  Se,  Te. 

Group  III — N,  P,  As,  Sb,  Bi. 

Group  IV — B. 

Group  V— Si,  C. 


Metals — 

Group  I — H. 

Group  II — K,  Rb,  Cs,  Na,  Li. 

Group  III— Ca,  Sr,  Ba. 

Group  IV— Be,  Mg,  Zn,  Cd,  Hg. 

Group  V — Pb,  Tl,  Cu,  Ag,  Au. 

Group  VI— Y,  La,  Cr,  Di,  Sr. 

Group  VII — Al,  Ga,  In. 

Group  VIII— Mn,  Fe,  Co,  Ni,  Cr,  Mo,  W,  U. 

Group  IX — Sn,  Ti,  Zr,  Th. 

Group  X— V,  Ta,  Nb. 

Group  XI — Pt,  Ru,  Rh,  Pd,  Ir,  Os. 

OXYGEN  (to  generate  acid). 

At.  Wt.  16.  Val.  II.  Symbol  0. 

Occurrence — 

% of  water,  Ys  of  air,  Y2  earth’s  crust,  % organized  matter. 

Preparation — 

By  heating  (1)  mercuric  oxide  HgO,  (2)  manganese  dioxide 
Mn02,  (3)  potassium  chlorate  KCIO3.  (4)  Decomposition  of 
water  (H2O)  by  electricity.  (5)  In  the  laboratory,  by  heating 
a mixture  of  KCIO3  + Mn02-  (6)  Commercially  obtained  by 

heating  barium  oxide  BaO  to  a red  heat  in  the  air.  This  forms 
barium  peroxide  Ba02,  which  heated  to  a white  heat  loses  half 
its  oxygen  and  is  reduced  again  to  the  oxide. 

BaC>2  — BaO  + 0. 

Properties — 

Colorless,  inodorous,  tasteless  gas — the  great  supporter  of 
combustion  and  life — combines  with  all  elements  save  fluorine 
to  form  oxides.  At  ordinary  temperatures  oxygen  is  mod- 
erately active,  but  its  chemical  affinity  increases  with  heat.  0 
is  slightly  soluble  in  water — with  low  temperature  and  great 
pressure  condenses  to  a liquid. 

Detection — 

Distinguished  from  all  gases  but  nitrous  oxide  N2O  by  its 
spark.  From  N2O  (1)  by  the  solubility'  of  0 in  potassium  pyro- 
gallate,  (2)  Sulphur  burns  in  0 but  not  in  N2O,  (3)  0,  exploded 
with  two  volumes  of  hydrogen  (H),  loses  all  its  volume,  while 
N2O  leaves  a residue  of  N. 


Combustion  — 

Rapid  oxidation  accompanied  by  light  and  heat. 

Decay — 

Slow  oxidation  of  organic  matter  assisted  by  bacteria. 

Kindling  Point — 

Temperature  at  which  bodies  combine  with  light  and  heat 
— Interposition  of  such  bodies  as  wire  gauze  serves  to  reduce 
the  temperature  below  the  kindling  point.  In  Davy  Safety 
Lamp,  a wire  gauze  jacket  surrounds  the  flame.  The  heated 
gases  passing  through  such  a jacket  are  cooled  below  the  kind- 
ling point,  and  there  can  be  no  combustion.  Small  quantities 
of  “fire  damp”  passing  through  the  gauze  from  without,  are 
heated  within  the  jacket  to  their  K.  P.  and  cause  a small  ex- 
plosion which  warns  of  a dangerous  proximity. 

Flame — 

A burning  gas  whose  luminosity  is  usually  due  to  red  hot  un- 
oxidized carbon  (C).  Intense  light  is  produced  by  the  glowing 
of  some  non-volatile  bodies  in  a flame  as  calcium  oxide  (CaO)  in 
calcium  light.  Other  things  being  equal,  combustion  produces 
less  light  and  more  heat  in  proportion  as  oxygen  is  increased. 
In  illuminating  tips,  gas  burns  by  oxygen  obtained  at  tip  of 
tube,  while  in  the  colorless  “Bunsen”  the  same  gas  is  well  mixed 
with  air  before  it  ignites.  In  the  oxv hydrogen  lamp  or  blow- 
pipe Oxygen  alone  is  mixed  with  Hydrogen  giving  an  intense 
non-luminous  heat.  The  comparative  heat  of  burning  bodies 
is  known  and  Carbon  (C)  yields  the  best  fuel  by  volume. 

Oxidation — 

Adding  of  0 to  any  substance. 

Reduction— 

Subtracting  of  0 from  any  substance. 

Reactions  of  Oxides — 

Oxides  may  have  one  of  three  reactions  to  litmus. 


Acid — 

Basic — 

Neutral — 

so2, 

k2o, 

h2o, 

no2, 

Na20, 

MgO, 

etc. 

etc. 

etc. 

Red. 

Blue. 

Purple. 

9 


OZONE,  03. 

Oxygen  occurs  in  two  alio  tropic  forms,  ordinary  oxygen 
O2  and  ozone  O3. 

Occurrence — 

Whenever  oxygen  is  prepared,  small  amounts  of  ozone 
form.  It  is  supposed  to  exist  in  the  air,  but  we  have  no  proof. 

Preparation — 

(1)  By  the  silent  discharge  of  an  electric  machine.  (2)  Slow 
oxidation  of  phosphorus.  (3)  In  the  laboratory  by  treating 
barium  peroxide  with  sulphuric  acid. 

BaC>2  + H2SO4  — 0(3)  + H2O  + BaS04. 

Properties  — 

“Condensed  oxygen”  is  a heavy  gas — poisonous — taste  and 
odor  like  weak  chlorine — at  ordinary  temperature  is  a powerful 
oxidizing,  bleaching  and  disinfecting  agent.  Ozone  acts  with  its 
extra  atom  of  0 — O3  = 02  + 0.  Here  0 does  the  work,  and 

02  (ordinary  oxygen)  is  set  free.  Ozone  condenses  to  a blue 
liquid — at  300°  breaks  to  O2. 

Detection — 

(1)  O3  blues  potassium  iodide  (KI)  starch  paper  (common 
to  Cl,  NO2  and  H2O2).  (2)  Blues  red  KI  litmus  paper.  This  is 

distinctive  in  the  absence  of  ammonia  (NH3).  Chief  use  of 

03  is  to  bleach  old  pictures. 

HYDROGEN-H. 

v. 

At.  Wt.  1.  Yal.  I. 

History — 

Discovered  by  Cavendish  in  1766. 

Occurrence — 

Free  in  volcanoes,  oil-wells,  and  from  decomposing  organic 
matter;  chiefly  combined  with  oxygen  in  H2O  and  organic 
bodies. 

Preparation — 

(1)  From  H2O  by  (a)  electricity,  (b)  metallic  sodium  or 


10 


potassium,  (c)  hot  iron.  (2)  Laboratory  method,  by  replacing 
H of  an  acid  by  a metal. 

Zn  + 2HC1  = H2  + ZnCl2. 

Properties — 

Tasteless,  inodorous,  colorless  gas  — does  not  support  com- 
bustion, but  burns  with  a blue  flame  to  H2O — mixed  with  oxy- 
gen explodes  at  K.  P. — lightest  of  gases,  H is  the  unit  of  specific 
gravity  for  gases — highly  diffusible  and  a strong  reducing  agent 
— H is  a gaseous  metal  playing  the  same  part  in  acids  as  the 
other  metals  play  in  salts— conducts  heat  and  electricity,  forms 
alloys,  and  in  electric  decompositions,  goes  to  the  negative 
(metal)  pole. 

Nascent  gas — 

Condition  at  immediate  generation  when  chemical  affinity 
is  much  stronger  than  usual — the  atoms  have  not  yet  united  to 
form  the  molecule. 

Nascent  atoms.  — Molecule. 

H + H = H2. 

WATER— H20. 

At  ordinary  temperature  a transparent  fluid  devoid  of 
taste  or  smell — thin  layers  colorless,  large  masses  blue — cooled 
to  a certain  temperature  water  c^stallizes,  forming  ice — this  is 
the  0°  C.— heated  to  a certain  temperature  water  boils — this  is 
the  100°  C.  On  cooling  water  contracts  till  4°  C.,  which  is  its 
maximum  density — below  4°  it  expands,  hence  ice  floats,  and 
water  pipes  break.  H20  evaporates  at  all  temperatures,  ice 
vaporizes  slowly  and  without  changing  to  liquid  form.  Heated 
water  changes  rapidly  to  an  invisible  gas  called  Steam,  which 
partially  recondenses  to  the  liquid  state  when  in  contact  with 
air. 

Composition  of  water — 

The  electric  current  resolves  water  into  two  volumes  of  H 
and  one  of  O ; conversely,  by  the  electric  current — or  by  heat — 
two  volumes  of  H and  one  volume  of  O always  combine  to  H20 
without  residue,  and  the  H20  thus  formed  may  be  heated  to 
two  volumes  of  steam.  Thus  water  by  volume  = H2  + 0,  or 
by  weight  = 2 H to  16  O,  or  1 H to  8 0. 


Distilled  water — 

Natural  waters  contain  many  substances  in  solution  and 
to  obtain  water  free  of  solids  it  must  be  boiled  and  the  steam 
condensed,  i.  e.,  distilled. 

Water  of  Crystallization — 

Molecules  of  water  which  enter  into  definite  combination 
with  many  salts,  when  these  crystallize  from  their  water  solu- 
tions. It  often  influences  the  color,  and  crystalline  form,  but 
may  be  driven  off  by  heat  without  alteration  of  substance. 

Efflorescence — 

Some  salts  yield  their  water  of  crystallization  to  the  air  as 
Na2C03,  10H26,  etc. 

Deliquescence — 

Some  salts  take  up  water  from  the  air  and  ultimately  melt, 
as  CaCl2,  NaNC>3,  etc. 

Unit  of  density — 

Water  at  15°  C.  is  unit  of  density  for  solids  and  liquids. 

Hydrometer — 

An  instrument  for  determining  specific  gravity  of  liquids 
with  reference  to  water  (1000) 


HYDROGEN  PER0XIDE-H202. 

History — 

Discovered  by  Thenard  in  1818. 

Occurrence — 

In  small  quantities  in  the  air  and  formed  wherever  ozone 
oxidizes  in  presence  of  H20. 

Preparation — 

When  barium  peroxide  Ba02  is  treated  with  H2S04  a 
double  reaction  takes  place.  Ozone  0(3)  is  generated,  which 
then  oxidizes  the  water  formed  to  H202. 

Ba02  + H2S04  = 0(3)  + H20  + BaSO*. 

03  + H20  = H202  + 02. 

Properties — 

“Ozonic  ether,”  “golden  fluid,”  “a  oxygenated  water,” 
H202  is  a thick  colorless  liquid.  Soluble  in  H20 — astringent 


taste — dilute  Cl  odor — strong  oxidizing  and  disinfecting  agent, 
contains  more  O than  any  other  substance  (94%), concentrated 
solution  easily  breaks  to  H2O  + O — dilute  slightly  acid  solution 
is  more  stable. 

Detection— 

(1)  Blues  KI  starch  paper. 

(2)  When  shaken  with  potassium  bichromate  (K^C^CL) 
and  ether  (CaHs^  0 it  gives  the  latter  a blue  color. 

Use- 

Antiseptic  spray,  and  to  bleach  old  pictures  and  hair. 
CHLORINE— Cl. 

At.  Wt.  35.5.  Val.  I. 

History — 

Discovered  by  Scheele  in  1774. 

Occurrence — 

Chiefly  as  common  salt  NaCl  in  sea  waters,  springs,  and 
especially  rock  salt  deposits.  Never  free,  because  of  strong 
affinities. 

Preparation — 

From  NaCl,  H2  SO4  and  Mn02  a triple  reaction. 

t (hydro- 

chloric) 

2NaCl  + H2SO4  = 2HC1  + Na2S04. 

(manganese  (manganese 

dioxide)  tetrachloride) 

Mn02  + 4HC1  — MnCL  + 2H20. 

MnCU  (heated)  = Cl2  + MnCl2 

breaking  up  of  MnCU  shown  by  change  of  color — gas  collected 
by  displacement,  or  over  hot  water  or  salt  solution. 

Properties — 

Green  \^ellow  gas — violent  odor,  poisonous;  strong  chemical 
agent  at  ordinary  temperature.  For  example,  Cl  extracts  H 
from  such  hydrocarbons  as  turpentine — with  some  metals  as 
antimony  Sb  it  combines  with  light  and  heat;  with  H it  unites 
in  direct  sunlight,  forming  HC1,  or  Cl  will  burn  in  H gas  to  form 
HC1  and  vice  versa.  Cl  is  a strong  disinfecting  agent  and  indi- 
rectly bleaches  through  the  action  of  nascent  0 which  it  frees 


from  H2O.  Cl  is  easily  soluble  in  H2O,  forming  a solution 
which  retains  the  properties  of  Cl  gas.  Cl  used  in  commerce 
chiefly  as  bleaching  powder. 

Bleaching  Powder  (see  Ca) — 

Is  formed  by  running  Cl  gas  over  slaked  lime  Ca(0H)2 — its 
approximate  composition  is  CaCl(OCl).  It  disinfects  and 
bleaches  chiefly  through  nascent  chlorine  which  is  freed  by  dilute 
acids. 

2CaOCl2  + 2HC1  m 2HC10  + CaCl2. 

HCIO  + HC1  — Cl2  + H20. 

Detection — 

By  its  odor,  bleaching  power,  and  by  turning  KI  starch 
paper  blue. 

CHLORINE  OXIDES. 

Chlorine  forms  a series  of  oxides  similar  to  that  of  nitrogen, 
2.  e.,  CI2O,  CIO2,  CI2O3,  CI2O5.  All  are  formed  indirectly  and 
unite  with  water  to  form  acids. 

CHLORIC  ACID— HCIO3. 

Preparation — 

The  anhydride  CI2O5  is  not  yet  known  in  the  free  state.  The 
K salt  of  the  acid  is  formed  by  passing  Cl  into  KOH, 

6KOH  + 6C1  — KCIO3  + 5KC1  + 3H20. 

The  K salt  is  then  treated  with  fluosilicic  acid  H2SiF6 
2KC10s  + H2SiF6  = 2HCIO3  + K2SiF6. 

Properties — 

A liquid  of  faint  odor;  strongly  acid  reaction;  easily  decom- 
posed ; a strong  oxadizing  and  bleaching  agent.  The  K salt 
is  a source  of  oxygen  supplying  enough  to  burn  some  com- 
bustible substances.  White  gunpowder  iscomposed  of  KCIO3 
and  sugar. 

HYPOCHLOROUS  ACID— HClO. 

The  anhydride  CI2O  is  a red  yellow  gas  with  a Cl  odor; 
explosive;  condenses  to  a liquid;  is  prepared  by  the  action  of  Cl 
on  HgO.  * 

mercuric  oxide 

HgO  + 4C1  = C120  + HgCl2. 


14- 


The  gas  dissolves  in  water  forming  HC10 
C120  + H20  = 2HC10. 

The  Ca  salt  Ca(C10)2  is  highly  important  being  the 
active  principle  of  “bleaching  powder.”  Hypochlorites  in 
general  are  very  unstable  and  act  as  strong  oxidizing  agents. 
The  other  oxides  of  Cl  are  unimportant. 

HYDROCHLORIC  ACID— H Cl. 

Occurrence — 

HC1  found  free,  in  the  gastric  juice  and  in  volcanoes — com- 
bined, chiefly  as  NaCl. 

Preparation — 

Generally  prepared  by  action  of  sulphuric  acid  on  common 
salt. 

H2SO4  + 2NaCl  = 2HC1  + Na2S04. 

May  be  formed  by  direct  union  of  H and  Cl  in  the  sunlight 
or  by  heat  or  electricity. 

Properties — 

Absolute  HC1  is  a colorless,  suffocating  gas — non-combust- 
ible, and  non-supporter — fumes  in  damp  air — is  eagerly  absorbed 
by  water,  forming  a 33  per  cent  solution  which  is  the  concen- 
trated HC1  of  the  laboratory — with  silver  nitrate  (AgNOs)  hy- 
drochloric acid  forms  a curdy,  white,  insoluble  precipitate  of 
silver  chloride,  AgCl,  which  is  the  test  for  HC1. 

BROMINE— Br. 

At.  Wt.  80.  Yal.  I. 

History — 

Discovered  by  Balard  in  mother  liquor  of  salt  works,  1827, 

Occurrence — 

Never  free — usually  as  K,  Na  or  Mg  bromide — in  salt  water 
(1  gr.  per  gal.),  rock  salt  deposits,  sea-weeds  and  springs — ac- 
companies chlorine  and  is  obtained  from  mother  liquor  of  salt 
works — Br  is  one  of  the  less  common  elements. 

Properties — 

Dark  red  liquid — red  brown  vapor — irritant  poison — char- 
acter similar  to  Cl  but  with  weaker  affinities  being  freed  from 


its  compounds  by  Cl — Chemical  unions  of  Br  often  accompanied 
by  light — Br  dissolves  in  water,  alcohol,  ether  and  chloroform. 

Preparation — 

Freed  from  its  salts  by  (1)  chlorine. 

2NaBr  + Cl2  = Br2  + 2NaCl. 

(2)  Manganese  dioxide  and  sulphuric  acid. 

2NaBr  + Mn02  + 2H2S04  = Br2  + K2S04  + MnS04  + 2H20. 

( Same  stages  in  this  reaction  as  in  preparation  of  Cl. ) 
Detection — 

Add  Cl  water  and  CS2  = red  brown  color  to  CS2.  Br  is  set 
free  by  Cl — CS2  dissolves  free  Br  with  a red  brown  color.  Br 
gives  starch  an  orange  color. 

Use — 

Br  is  used  in  dye  factories,  photography,  medicine. 


HYDROBROMIC  ACID— HBr. 

Properties — 

Colorless  gas — fumes  in  moist  air,  eagerly  absorbed  by  water, 
forming  HBr  2H2O.  Analogous  to  HC1,  but  less  stable,  is  de- 
composed by  H2SO4.  Bromides  freed  by  chlorine. 

Preparation — 

(1)  Direct  union  of  H and  Br  at  red  heat. 

(2)  Generally  prepared  by  action  of  phosphorus  tri bromide 
and  water. 

(Phosphorous  acid) 

PBr3  + 3H20  = H3PO3  + 3HBr 

(3)  Cannot  be  made  pure  from  salt  by  H2SO4,  as  resulting 
HBr  is  decomposed  bjr  H2SO4. 

Solubility — 

Most  Bromides  are  soluble  in  H2O 

OXIDE  AND  CHLORIDE  OF  Br. 

Bromine  chloride,  BrCl — 

Br  absorbs  Cl  forming  a yellow  unstable  liquid  with  bleach- 
ing properties.  Above  10°  C.  changes  to  Br  and  Cl. 


Bromine  oxides — 

Bromine  and  oxygen  unite  with  difficulty  to  form  a series 
entirely  analogous  to  Cl  and  made  in  same  way — 

Br20,  — , Br02,  Br203,  Br20s. 

IODINE— I. 

At.  Wt.  127— Val.  I. 

History — 

Discovered  by  Courtois  in  1811. 

Occurrence — 

Not  free — compounds  accompanying  Cl  and  Br,  in  sea-water, 
springs,  etc. — Chief  source  is  the  ash  of  sea-weeds  called  “kelp.” 

Prope  rties — 

Grey  black  solid  with  violet  vapor — weak  Cl  odor — difficultly 
soluble  in  water — dissolves  easily  in  alcohol  or  a solution  of  KI — 
chemically  similar  to  Cl  and  Br  but  with  weaker  affinities,  is 
freed  by  either  from  its  compounds. 

Preparation- 

Kelp  solution  is  partly  evaporated — I remains  in  the  mother 
liquor— solution  is  then  distilled  with  MnC>2  and  H2SO4. 

2KI  + Mn02  + 2H2S04  = I3  + M11SO4  + K2S04  + H20  ' 

(Game  reaction  as  for  chlorine). 

Detection — 

(1)  With  Cl  water  and  CS2  gives  violet-color. 

(2)  With  cold  starch  paste,  a blue  color. 

Used- 

In  dye  factories  and  photography. 

HYDRIODIC  ACID — HI. 

Properties— 

Colorless  gas  with  suffocating  odors — fumes  in  moist  air — 
absorbed  by  water— easily  decomposed;  hence,  cannot  be  made 
with  H2SO4— Strong  reducing  agent. 

Preparation — 

(1)  From  Phos-tri-iodide  and  water. 

PI3  + 3H20  = H3PO3  + 3HI  (strong  acid). 

(2)  I with  H2S  gives  weak  acid. 

I2  + h2S  = S + 2HI. 


17 

OTHER  COMPOUNDS  OF  IODINE. 


Iodine  Chlorides  (IC1,  IC18) — 

Two  compounds  IC1,  ICI3  formed  by  direct  action  of  the 
elements;  used  to  add  I to  organic  bodies. 

Iodine  Bromide  (IBr) — 

IBr  is  a solid  similar  to  I and  formed  as  IC1. 

Iodine  Oxides — 

Iodine  forms  the  same  oxides  as  Br  or  Cl;  I2O,  IO2, 
I2O3,  l20£. 


Iodic  Acid  (HI03)— 

A white  solid  formed  by  boiling  I in  HNO3;  permanent  in 
air,  but  with  heat  changes  to  I2O5  +H2O;  forms  salts  similar  to 
chlorates. 


FLUORINE— FI. 

At.  Wt.19.  Val.  1. 

History — 

Discovered  by  Davy  in  1812,  but  not  isolated  till  1886  by 
Moissan. 

Occurrence — 

Chiefly  in  fluor  spar  or  fluorite,  CaF2,  and  cryolite 
AlFla,  3NaFl. 

Properties — 

Colorless  gas — deadly  poisdn — attacks  most  substances,  but 
does  not  combine  with  O,  C,  Pb  or  Pt,  nor  attacks  guttapercha — 
properties  analogous  to, chlorine  but  much  stronger,  its  atomic 
weight  being  much  less.  F frees  Cl,  Br  or  I from  their  com- 
pounds. ’• 

HYDROFLlfbRIC  ACID,  HF. 

Preparation 

in  Et  retort 


•%  *»i 

From  fluor  sparjfov  j 

daft  + Ullft  ^aS04  + 2HP. 

Properties—  , . L 

N Colorless,  mobile  liquid — fumfes  in  air — corrosive  poison — 
dissolves  all  metals  ex^pi  J^b;:  Pt  and  Au,  and  all  oxides,  in- 

• *' 

1 


* "' -*'**■ 


IS 


eluding  SiC>2.  Must  be  kept  in  gutta  percha  bottles.  Used  to 
etch  glass  and  detected  by  its  etching  quality. 

SULPHUR— S. 

At.  Wt.  32.  Val.II. 

History — 

Used  by  the  ancients  for  fumigation  and  medicine — Consid- 
ered by  alchemists  to  be  the  principle  of  combustion. 

Occurrence — 

(1)  Native  in  volcanic  countries,  especially  Sicily. 

(2)  Sulphates  and  Sulphides. 

(3)  Organic  compounds. 

The  amount  of  S produced  is  about  375,000  tons,  of  which 
nine-tenths  comes  from  Sicily. 

Extraction  from  Ores — 

(1)  Melting  ores  in  a pot,  S floats  and  is  dipped  out. 

(2)  Burning  ores  in  small  supply  of  air  so  that  most  S 
melts. 

Purified — 

By  distillation — That  which  passes  over  first  cools  more 
rapidly  forming  flowers  and  is  less  pure  than  the  roll  which 
forms  afterwards — Flowers  contain  H2SO4;  H2S  and  SO2  as 
impurities. 

Properties— 

Strong  affinities — Union  with  metals  often  produces  light 
and  heat.  Moist  Fe  -f-  S = FeS  + heat — unites  with  haloids  at 
ordinary  temperature — Slightly  heated  S + P explodes — H run 
into  melted  S = H2S;  Compounds  of  S alone  are  called  sul- 
phides. S is  wholly  analagous  to  O and  when  oxides  of  an 
element  are  soluble,  usually  sulphides  are.  With  O,  S burns 
with  blue  flame  to  SO2. 

Uses — 

Disinfectant;  making  H2SO4,  SO2,  matches  and  gunpowder. 

HYDROGEN  AND  SULPHUR. 

Two  compounds  well  known— hydrogen  sulphide,  H2S, 
and  persulphide,  H2S2.  The  latter  is  unimportant. 


Hydrogen  Sulphide  (H2  S)— Occurrence- 

Free  in  mineral  springs,  and  when  organic  bodies  decay. 

Preparation — 

(1)  From  a sulphide  and  an  acid — 

FeS  + 2HC1  = H2S  + FeCl2. 

(2)  Chemically  pure — 

CaS  + 2HC1  = H2S  + CaCl2 
Collect  over  warm  water  or  salt  solution. 

(3)  Formed  by  heating  elements  to  400°  C. 

Properties— 

Colorless  gas — bad  odor — poisonous — combustible,  burns  to 
S02  and  H2O— mixture  with  air  explodes— unstable,  oxidizes  in 
air  at  ordinary  temperature  to  S + H20.  Easily  unites  with 
metals  to  form  sulphides.  Dissolves  in  H20  (3  volumes  to  1), 
and  water  solution  has  properties  of  the  gas,  decomposing  in 
light  to  H20  + S.  HoS  is  a dibasic  acid,  hence  forms  two  series 
of  Salts,  viz.,  KHS  =.  potassium  sulphydrate, 

K2S  = potassium  sulphide. 

SULPHUR  HALOIDS. 

S unites  with  Cl  in  three  proportions— S2C12  (SCI),  SC12, 
SCI4.  The  monochloride  S2C12  is  a red  yellow  liquid  formed  by 
passing  Cl  over  molten  S — used  in  vulcanizing  rubber;  SCl2and 
SCU  are  highly  unstable. 

With  Br  and  I,  S forms  similar  compounds. 

SULPHUR  DIOXIDE,  S02. 

Occurrence — 

Native  in  volcanic  gases. 

Preparation — 

(1)  Oxidizing  sulphur — 

S + 02=r  S02. 

(2)  Reducing  sulphuric  acid  (H2S04)— 

Cu  + 2H2S04  = CuS04  + S02  + H20. 

S02  is  called  “Sulphurous  acid”  because  the  acid  proper, 
H2S03,  is  unstable. 


Properties — 

Colorless*  suffocating  gas — inhibits  combustion — poisonous 
to  plants  and  animals — disinfecting  agent — bleaches,  but  color 
is  restored  by  alkalies — easily  condensed  without  pressure — 
highly  soluble  in  H20  (50  volumes  to  1)— salts  easily  oxidize 
to  sulphates,  hence  good  reducing  agents.  H2SO3  is  a weak, 
dibasic  acid,  little  stronger  than  carbonic— S02  does  not  unite 
directly  with  O but  with  ozone  forms  SO3. 

SULPHUR  TRIOXIDE— S03. 

Preparation — 

Sulphuric  anhydride  formed  from  S02  by  (1)  ozone,  (2)  O 
over  divided  Pt,  (3)  electric  spark,  (4)  distilling  disulphuric 
acid,  H2S2O7. 

Properties — 

White,  silky,  crystalline  needles — fumes  in  damp  air— acid 
only  with  H20.  Decomposes  at  high  temperature  to  S02  + O. 
Unites  with  H20  with  great  heat. 

SULPHURIC  ACID-H2SO4. 

History — 

Known  for  centuries — alchemists  prepared  it  from  S and 
nitric  acid  (HNO3)  or  by  heating  FeS04,  calling  it  “oil  of  vitriol” 
— first  lead  chambers  built  in  1746. 

Occurrence — 

Free  in  volcanic  rivers  and  the  fluids  of  some  mollusca — 
combined  as  sulphates  in  large  quantity  as  CaS04,  etc. 

Preparation — 

(1)  Oxidation  of  S by  HNO3  directly. 

(2)  Oxidation  of  S02  by  HNO3  in  presence  of  H20. — S02 
from  burning  FeS2  (iron  sulphide),  with  O from  the  air,  is 
passed  into  lead  chambers,  through  which  HNO3  and  H20  are 
streaming — HNO3  is  decomposed  to  the  oxides  of  N — these 
yield  up  O to  S02  forming  SO3,  which  unites  with  H20  to  form 
H2SO4.  The  lower  oxides  of  N reunite  with  O of  the  air, 
forming  the  higher,  N02  and  N2O3,  which  again  oxidize  a new 
portion  of  S02. 

The  dilute  chamber  acid  (sp.  gr.  1.5(c)  is  concentrated  in 


Pt  or  glass  vessels  to  commercial  strength  (1.830).  Crude 
H0SO4  contains  as  impurities  PbSCU,  HNO3,  often  As  from  S 
or  FeS2  used.  Impurities  are  generally  removed  by  distillation, 
but  an  arsenic-free  acid  is  obtained  only  by  using  arsenic-free 
materials. 

* 

Properties— 

Heavy,  svrup-like  liquid,  colorless  when  pure,  usually 
brown  from  charred  organic  matter — odorless — corrosive — in- 
tensely acid.  Heated  above  boiling,  gradually  breaks  up  to 
PI2O  + SO3 — mixed  with  water  great  heat  is  evolved,  due  to 
formation  of  higher  hydrates — H2SO4  + H2O  or  2H2O.  Sul- 
phates of  strong  bases  are  not  easily  broken  up  by  heat. 

FUMING  OR  NORDHAUSEN  ACID— H2S207. 

By  running  SO3  into  concentrated  H2SO4  we  get  H2S2O7, 
disulphuric  acid,  a white  solid  stronger  than  H2SO4  and  used 
in  indigo  and  alizarin  manufacture. 

SELENIUM  (Se)  AND  TELLURIUM  (Te). 

At.  Wt.  79.  At.  Wt.  128. 

Two  rare  elements  closely  analogous  to  Sulphur — Tellu- 
rium was  discovered  in  1782  in  ores  of  gold — physical  pro- 
perties are  metallic,  chemically  it  acts  like  S. — Selenium  discov- 
ered in  1817  by  Berzelius  from  the  mud  on  floor  of  lead  cham- 
bers. HNO3  with  Te  or  Se  forms  respectively  tellurious  acid, 
H2Te03  or  selenious  acid,  H2Se03 — with  KNO3  the  elements 
form  K2Te04  or  K2SeC>4.  Selenium  colors  the  flame  blue  and 
H2Se  is  remarkable  for  its  odor.  Selenium  changes  its  resis- 
tance to  an  electric  current  when  exposed  to  light. 

SULPHUR  GROUP. 

0 16— S 32— Se  79.4— Te  128. 

O a gas — others  solids — O and  S are  non-metals — Se  partly 
metallic — Te  a metal.  All  unite  with  H2 — the  three  solids  form 
with  H2  volatile  strong  smelling  gases  of  acid  nature — Stabil- 
ity of  H compounds,  varies  inversely  as  atomic  weights,  of  O 
compounds  directly — O is  always  bivalent,  the  others  vary  in 
valence 


22 


NITROGEN— N. 

At.  Wt.  14.  Val.  III.  or  V. 

History — 

Discovered  in  1772  by  Rutherford — Scheele  and  Lavoisier 
showed  that  air  = O + N,  called  azote  by  the  French  and  Ital- 
ians. « 

Occurrence — 

% of  the  air  and  in  many  organic  substances. 

Preparation — 

(1)  Removing  O from  air  by  (a)  Phosphorus — (b)  red-hot 
Cu. 

(2)  Heating  Ammonium  nitrite  (H4N)N02 

(H4N)N02  = N2  + 2H20. 

Properties — 

Characterized  by  its  inertness  at  ordinary  temperature — 
colorless — inodorous,  tasteless  gas — non-combustible — non-sup- 
porter— not  poisonous — at  white-heat  unites  with  metals— many 
of  its  compounds  are  explosives,  while  all  explosives  contain  N. 

AIR. 

A comparatively  constant  mixture  of  % N,  % 0,  with 
appreciable  quantities  of  H2O  and  CO2  and  trace  of  NH3  com- 
pounds, NaCl,  and  dust.  N and  0 are  quite  constant;  the  oth- 
ers vary. 

Impurity  of  air  is  usually  determined  by  amount  of  CO2 
present.  CO2  of  itself  is  not  harmful  in  small  quantity,  but  ex- 
perience has  shown  that  it  bears  a constant  ratio  to  the  amount 
of  ammonia  compounds  present.  This  “organic  waste,”  so- 
called,  is  intensely  poisonous,  although  it  exists  in  quantity  too 
small  for  easy  determination. 

Air  has  the  same  physical  properties  as  the  gases  of  which 
it  is  composed,  being  a mixture,  not  a compound,  because  (a) 
it  varies  somewhat  in  composition,  (b)  proportion  of  O and  N 
is  not  governed  by  atomic  weight,  (c)  absorbed  by  water  as 
free  gases  35  vol-  O : 65  vol.  N,  not  1 vol.  0 : 4 vol.  N. 

AMMONIA— NH3. 

History — 

Called  “spirits  of  hartshorn”  because  formerly  obtained 


23 


from  horns  of  deer.  Gas  discovered  by  Priestley,  who  called  it 
alkaline  air. 

Occurrence — 

Not  free  in  nature  as  NH3.  Found  as  salt  in  animal  secre- 
tions, especially  urine.  Found  also  when  nitrogenous  organic 
bodies  decay  or  are  distilled,  hence  found  in  liquor  of  gasworks, 
as  coal  contains  N compounds. 

Preparation — 

Commercially  obtained  from  gas  works.  In  laboratory 
from  ammonium  chloride  (H4N)  Cl  and  calcium  hydrate  Ca 
(0H)2.  x 

2H4NCI  -I-  Ca(OH)2  = 2NH3  + CaCl2  + 2H20. 

Properties — 

Colorless  gas — pungent  odor — strong  alkali — non-supporter 
of  combustion — burns  feebly  in  the  air.  Decomposed  by  electric- 
ity or  by  chlorine  in  sunlight. 

4NH3  + 3C1  — 3H4NC1  + N. 

NH3  is  highly  soluble  in  H20,  forming  H4NOH,  ammonium 
hydrate,  a strong  base  in  which  the  radicle,  H4N,  deports  itself 
like  the  metals  Na  or  K,  and  will  be  treated  with  them. 

NITROGEN  HALOIDS 

Do  not  form  directly  from  elements — all  are  highly  explosive 
— most  important  is  NCI3 — a dark  red  liquid,  formed  by  action 
of  Cl  on  H4NCI. 

H4NCI  + 6C1  = NCls  + 4HC1. 

NBr3  and  NI3  are  similar. 

NITROGEN  OXIDES. 

Five  known  compounds — 

Nitrous  oxide.  Nitrous  anhydride.  Nitric  oxide.  Nitrogen  peroxide.  Nitric  anhydride 

N20.  N203.  N0(N202)  N02(N204).  n2o5. 

The  first,  second  and  fifth  act  as  anhydrides,  forming  hypo- 
nitrous,  nitrous,  nitric  acids.  The  third  and  fourth  form  a mix- 
ture of  HN02  and  HNO3.  All  the  oxides  are  prepared  from 
HN03. 

Nitric  Anhydride  N20s — 

Discovered  by  Deville.  Made  by  passing  Cl  over  dry 
AgN03— apparatus  entirely  of  glass,  melted  together. 


24 


2AgN03  + Cl2  = 2AgCl  + N205  + O. 
White,  unstable  solid — with  H2O  forms  HNO3. 

N205  + H20  = 2HN03. 


NITRIC  ACID— HN03. 

History — 

Known  early  as  9th  century  when  it  was  made  by  distilla- 
tion of  ZnSC>4  + KNO3  and  called  ‘‘Aqua  fortis,”  “Parting 
water,”  or  “Spiritus  nitri  fumans  Glauberi.” 

Occurrence — 

Not  free — in  salts  as  nitrates,  is  widely  distributed  in  the 
earth  especially  in  Chili  as  NaN(>3  or  Chili  Saltpeter — in  small 
quantities  in  air  and  water. 

Formation — 

From  elements  by  electric  spark  in  presence  of  H2O — also  in 
process  of  decay,  occasioned  by  bacteria,  the  so-called  “nitrifi- 
cation,” which  takes  place  best  in  the  dark  and  is  stopped  by 
destruction  of  the  bacteria  through  chloroform  or  boiling. 

Preparation — 

Action  of  H2SO4  on  K salt. 

2KN03  + H2S04  — 2HN0s  + K2S04. 

Impurities — 

Ordinary  KNO3  contains  KC1;  hence,  HC1  is  formed  with 
crude  HNO3 — another  common  impurity  is  H2SO4. 

Properties — 

Pure  HNO3  is  a colorless  volatile  liquid  leaving  no  residue- 
no  stronger  acid  has  been  yet  obtained — so  powerful  an  oxidiz- 
ing agent  that  it  deflagrates  with  easily  combustible  substances 
— exposed  to  sunlight  it  gradually  decomposes  to  its  oxides  and 
H2O;  hence,  strong  acid  is  often  colored  by  NO2 — a powerful 
mono-basic  acid;  it  dissolves  most  of  the  metals  forming  nitrates 
all  of  which  are  soluble  in  water,  some  insoluble  in  HNO3,  all 
decompose  at  high  heat. 

Used— 

Commercially  to  etch  copper. 


25 


Aqua  Regia — 

Formed  by  mixing  HNO3  + HC1;  acts  as  nascent  chlorine, 
forming  a chloride. 

2HN03  + 6HC1  = 4H20  + 2N0C1  + Cl4. 

A long  known  mixture  of  dark  yellow  color  and  suffocating 
fumes  and  odor — called  “royal  water”  because  it  dissolves  “no- 
ble” metals,  gold  and  platinum. 

Nitrogen  Peroxide  NO2 — 

Whenever  NO  is  generated  it  combines  with  0 of  air  to  NO2 
— red  brown  gas — poisonous,  suffocating  odor — very  corrosive — 
colors  organic  bodies  yellow — strong  oxidizing  agent,  C and  P 
burn  in  its  vapor — easily  liquified  and  solidified.  NO2  does  not 
form  a hydrate,  but  with  H2O  gives 

2N0S  + H20  = HN02  + HNOs. 

Nitrous  acid  HNO2 — 

Hydrate  of  N2O3  and  formed  by  heating  HNO3  with  starch, 
sugar,  or  other  easily  oxidizable  substance,  which  reduces  HNO3. 
In  same  way  nitrites  are  prepared  by  reduction  of  nitrates — 
KNO3  + Pb  = PbO  + KN02. 

Nitrites  are  mostly  soluble  in  H2O. 

Nitrous  anhydride  N2O3 — 

Formed  by  decomposition  of  nitrites — a red  brown  gas 
easily  condensed  to  indigo  blue  liquid.  Passing  the  gas  into 
alkali  hydrates  forms  corresponding  nitrites — 

2K0H  + N203  ==  2KN02  + H20. 

Nitric  Oxide  NO — 

Never  found  free,  as  it  combines  with  0 of  air  to  form  NO2. 
Made  by  action  of  HNO3  on  Cu. 

3Cu  + 8HNO3  = 3Cu(N03)2  + 2N0  + 4H20. 

At  high  temperature  N2O  is  formed. 

Colorless  gas,  irrespirable,  poisonous — oxidizes  in  air  to 
NO2 — slightly  soluble  in  H2O.  NO  supports  combustion  of  P 
and  S but  not  of  ordinary  substances,  as  wood. 

Myponitrous  Acid  HNO — 

Discovered  in  1871  by  reducing  KNO3  with  NaHg.  Known 
only  as  salt. 


Nitrous  oxide  NoO — 

Discovered  by  Priestley  1776. 

Colorless  gas — sweet  taste — will  not  support  life,  but  can 
be  breathed  a short  time,  producing  intoxication  and  anaes- 
thesia (laughing  gas).  N2O  supports  combustion  like  O but 
extinguishes  burning  S.  Condenses  easily  and  forms  with 
ether  (C2Hs)20  the  coldest  known  mixture — distinguished 
from  O by  leaving  a residue  of  N when  exploded  with  H. 

Prepared  by  heating  (H4N)NC>3  (ammonium  nitrate) — 
(H4N)N03  = n20  + 2H20. 

At  high  temperature  NO  forms  with  explosion. 


Valence 

is  that  property  of  an  element  by  which  it  combines  with  defi- 
nite masses  of  other  elements.  Unit  of  valence  is  hydrogen. 
Elements  are  classified  according  to  their  power  of  uniting  with 
or  replacing  different  proportions  of  the  unit.  Thus  chlorine 
unites  with  one  H to  form  HC1  and  is  a univalent  element. 
Oxygen  forms  H2O,  requiring  two  parts  of  H and  is  a Divalent 
element.  Nitrogen  forms  H3N,  acting  as  a trivalent  element. 
Carbon  forms  H4C,  acting  as  a quadrivalent  element.  In  a 
comparatively  few  instances  valence  requires  five,  six  or  seven 
InTlrogens.  Under  the  same  conditions,  valence  of  a given  element 
is  always  constant,  but  it  may  vary  for  the  same  substance  in 
different  compounds,  depending  on  what  element  it  unites  with 
and  under  what  conditions  union  takes  place.  Generally  when 
a substance  is  present  in  smaller  mass  it  unites  with  higher 
valence.  Thus  a small  amount  of  carbon  heated  in  presence  of 
much  oxygen  forms  CO2,  where  C is  quadrivalent,  but  where  a 
large  amount  of  C is  heated  with  small  0 the  gas  formed  is  CO, 
where  C acts  as  a bivalent  element.  The  greatest  variation 
appears  in  the  oxides  and  chlorides.  The  oxides  of  N for  exam- 
ple— N20,N202,N203,N204,N205— give  to  nitrogen  five  distinct 
valences ; and  other  elements  form  similar  series. 

In  general,  valence  may  be  considered  less  as  a property  of 
the  specific  elements  and  more  as  a function  incident  to  their 
combination. 


PHOSPHORUS— P. 

At.  Wt.  31.  Val.  Ill  or  V. 

History — 

Discovered  in  1667  by  Brand  when  searching  for  philoso- 
pher’s stone  in  urine — which  was  the  only  source  till  1750 — 
later  Scheele  obtained  it  from  bones. 

Occurrence — 

In  soil  as  apatite  and  pyromorphite — extracted  from  soil 
by  plants — then  to  animals  in  bones,  brain  and  urine. 

Properties — 

Occurs  in  two  important  allotropic  forms,  yellow  crystal- 
line and  red  amorphous  phosphorus.  Yellow  is  converted  by 
heat  (250°)  to  red — ordinary  P exposed  to  light  becomes  red 
on  surface — common  or  yellow  P is  a wax-like  solid — ozone 
odor — luminous  in  the  dark — oxidizes  easily— burns  at  44° — 
a mixture  of  P and  KCIO3  detonates  when  struck — causes 
chronic  poisoning — almost  insoluble  in  water,  but  dissolves 
easily  in  CS2 — preserved  in  H2O. . Red  P is  a weaker  chemical 
agent — odorless,  non-phosphorescent — not  poisonous — does  not 
dissolve  in  CS2 — is  stable  at  ordinary  temperatures — at  260°  C 
red  changes  back  to  yellow. 

Preparation — 

Bones  are  burned  to  remove  organic  matter  ( 55%  whole). 
Ash  treated  with  H2SO4 — this  changes  insoluble  Ca3(P04)2  of 
the  bones  to  a soluble  phosphate  which  heated  with  charcoal 
gives  free  P — the  essential  reaction  with  Cis  2P2O5  + 5C  = 
P4  + 5CO2.  P is  purified  by  redistillation.  Cannot  be  success- 
fully prepared  on  small  scale. 

Use— 

Chiefly  for  making  matches.  Safety  matches  = mixture  of 
potassium-chlorate  (KCIO3)  and  antimony  trisulphide  (Sb2Ss) 
— the  surface  = red  phosphorus  + Mn(>2.  Ordinal  match  = 
yellow  P,  some  oxidizing  agent  as  MnC>2  + KCIO3  and  glue. 

PHOSPHINE— PH3. 

Properties — 

The  only  important  compound  of  P and  H — a gas  of  disa- 


greeable  odor  and  highly  poisonous — chemical  action  similar  to 
NH3— forms  phosphonium  compounds  analogous  to  H4N  com- 
pounds. P and  H form  also  P2H4,  a liquid,  and  P4H2,  a gas, 
but  these  are  unimportant.  PH3  does  not  ignite  spontaneous- 
ly, except  in  presence  of  P2H4. 

Preparation — 

From  KOH,  P and  H2O  forming  PH3  and  an-  acid  of  P. 

3 KOH  + P4  + 3 H20  = PH3  + 3 KH2PO2. 

Phosphorus  trichloride  PCI3 — 

The  more  important  chloride  of  P — prepared  by  action  of 
Cl  on  melted  P — a colorless  liquid  which  fumes  in  moist  air  and 
with  water  forms  phosphorus  acid. 

2 PC13  + 6 H20  = 6 HC1  + 2 H3PO3. 

There  is  also  a penta-chloride  prepared  in  same  way  but  with 
excess  of  Cl — a white  crystalline  substance. — I and  Br  form 
similar  compounds,  all  of  which  are  much  used  in  the  arts. 

Phosphorus  Oxides — 

P forms  two  oxides  by  burning  in  greater  or  less  quantity 
of  dry  air,  P2O3  and  PO5 — both  white  feathery  solids — P2O5  is 
more  important  ancr  generally  called  phosphoric  ant^dride 
— unites  eagerly  with  water  in  different  proportions  to  form 
acids. 

(1)  P2O6  + H2O  — 2HPO3  (Meta- or  glacial  phosphoric 

acid). 

(2)  P2O5  + 2H2O  = H4P2O7  (Pyrophosphoric  acid). 

(3)  P2O5  + 3H2O  = 2H3PO4  (Orthophosphoric  acid). 

Orthophosphoric  acid  H3PO4 — 

P is  oxidized  with  HNO3  and  residue  evaporated — difficult 
operation  as  oxidation  is  slow  and  does  not  work  with  weak 
while  it  explodes  with  strong  HNO3 — can  be  made  also ‘from 
the  Ca  salt  but  is  difficult  to  purify. 

Salts  — 

Form  three  classes — M2PO4,  M2HPO4,  M3PO4.  Most  sta- 
ble are  the  alkaline  M2HPO4  or  of  heavy  metals  M3PO4. 

By  ignition,  M3PO4  unchanged. 

B3’  ignition,  M2HPO4  changed  to  pyrophosphate  M4P2O7. 
By  ignition,  MH2PO4  changed  to  metaphosphate  MPO3. 


Pyrophosphoric  Acid  H4P2O7— 

Prepared  by  heating  H3PO4  or  adding  H2O  to  P2O5  as 
above. 

Gradually  absorbs  H2O  forming  ortho-  acid. 

H4P2O7  is  a tetrabasic  acid  but  acts  as  dibasic,  forming 
M4P2O7  and  M2H2P2O7. 

All  normal  salts  are  soluble  save  those  of  alkalies. 

All  ortho-  salts  change  with  heat  to  pyro-  salts. 

Metaphosphoric  Acid  HPO3 — 

Ortho-  acid  slightly  ignited  gives  pyro-  and  strongly  gives 
meta- — the  melted  mass  solidifies  to  “glacial”  meta-  acid — this 
dissolves  in  H^Oand  acts  like  P2O5  in  H2O— solution  coagulates 
albumen  and  gives  white  precipitate  with  AgN03  or  BaCl2 — it 
is  gradually  changed  to  ortho-. 

Phosphorous  Acid  H3PO3 — 

1)  . PC13  + 3H20  = H3PO3  + 3HC1. 

2)  Slow  oxidation  of  P in  moist  air. 

P2O3  + 3H20  = 2H3PO3. 

Strong  reducing  agent,  oxidizing  to  H3PO4. 

Forms  two  salts,  MH2PO3  and  M2HPO3. 

Hypophosphorous  Acid  H3PO2— 

Anhydrite  P2O  has  not  been  obtained— free  acid  formed  by 
treating  Ba  salt  with  H2SO4 — Ba  salt  formed  in  same  reaction 
as  for  PH3. 

3Ba(0H)2  + 8P  + 6H20  = 3 Ba(H2P02)f,  + 2PH3. 

White  crystalline  mass — a monobasic  acid — Ca  salt  used  in 
medicine. 

ARSENIC— As. 

At.  Wt.  75.  Yal.  I,  III  or  V. 

History — 

Element  long  known. 

Occurrence — 

Widely  distributed  as  native  As  and  compounds,  especially 
sulphides  AS2S2  realgar,  AS2S3  orpiment  and  mispickel  FeAsS— 
As  known  in  commerce  as  “cobalt”  or  “fly stone.” 


30 


Properties — 

Two  allotropic  modification — amorphous  black  powder, 
and  the  common  crystalline  form— a brittle,  steel  grey  solid 
N with  metallic  lustre— vapor  is  yellow  with  garlic  odor— burns 
with  blue  flame  to  AS2O3 — metallic  As  is  not  poisonous. 

Preparation — 

(1)  Heating  mispickel  in  clay  cylinders.  \ 

(2)  Heating  AS2O3  with  charcoal  and  resubliming  the  pro- 
duct. 

Use- 

Pigments— flypaper — also  forms  alloys  as  with  Pb  to  make 
shot. 

HYDROGEN  ARSENIDE,  ARSINE— AsH3. 

history — 

Discovered  by  Scheele. 

Properties — 

Colorless  gas — disagreeable  odor — violent  poison — burns 
with  blue  flame  to  AS2O3  + H2O — with  limited  supply  of  air  to 
As  + H2O — Decomposed  at  red  heat  to  H + As  which  is  deposited 
as  a mirror  giving  a delicate  test  for  As. 

As  in  reagents — FeS2  usually  contains  some  arsenic,  whence 
it  passes  into  H2SO4,  HC1,  and  any  reagents  made  from  these 
acids. 

Preparation — 

(1)  By  heating  the  alloy  As2Zn3  with  an  acid. 

AS2Ztl3  + 3H2SO4  = 2Astt3  + Z11SO4. 

(2)  Formed  when  As  is  treated  with  nascent  H. 

As  + H3  = AsH3. 

(3)  When  compounds  of  As  are  treated  with  organic  bodies 
especially  decaying  bodies,  thus  ASH3  is  formed  by  As  in  wall 
paper. 

Arsenic  trichloride  (ASCI3) — 

Formed  by 

(a)  burning  As  in  Cl. 

(b)  As203  6HC1  = 2AsC13  + 3H20. 

— Heavy,  colorless,  fuming  liquid— important  from  its  vola- 
tility as  it  may  cause  loss  of  As  in  analysis. 


31 


As — OXIDES — ACIDS  AND  SULPHIDES. 

Oxides — 

The  oxides  AS2O3  and  AS2O5  correspond  to  those  of  P and 
N and  are  of  acid  nature — AS2O5  combines  directly  with  water, 
forming  H3ASO4,  while  AS2O3  does  not  easily  dissolve;  hence  is 
known  as  arsenious  acid . 

Arsenic  trioxide  (AS2O3) — 

The  common  compound  of  As — commercially  known  as 
“arsenic”  or  “white  arsenic” — a white  amorphous  powder — 
formed  when  As  burns — slightly  soluble  in  H2O  with  faint  acid 
reaction — odorless — sweet  metallic  taste — deadly  poison — AS2O3 
unites  with  bases,  forming  arsenites — H2S  precipitates  AS2S3, 
which  is  soluble  in  (H4N)2S. 

Arsenic  pentoxide  (AS2O5) — 

A colorless  deliquescent  mass — formed  when  As2C>3is  heated 
to  a red  heat— not  when  As  burns— its  water  solution  forms 
arsenic  acid  (H3ASO4.) 

Arsenic  Acid  (H3ASO4) — 

White  crystalline  solid,  formed  when  AS2O3  is  digested  in 
HN03. 

(ortho-arsenic  acid) 

As203  + 2HN03  + 2H20  = 2H3As04  + N2O3. 

If  heat  is  applied  the  pyro-  and  metarsenic  acids  are  formed 
corresponding  in  formula  to  the  phosphoric  series — Cu3(As(>4)2 
is  blue — Ag3As04  brown — with  H2S  arsenic  acid  is  generally 
first  reduced  to  AS2O3  then  precipitated  as  AS2S3. 

D 2H3ASO4  + 2H2S  = AS2O3  + 5H20  + 2S. 

As203  + 3H2S  = As2S3  + 3H20. 

ARSENIC  SULPHIDE. 

Chief  is  Orpiment  AS2S3 — volatile  without  decomposition — 
heated  with  reducing  agent  to  As — soluble  in  alkaline  sulphides, 
forming  salts — 

As2S3  + 3(H4N)2S  = 2(H4N)3AsS8. 

A less  important  sulphide  is  Realgar  (AS2S2). 


ANTIMONY— Sb. 

At.  Wt.  120.  Val.  I.  III.  or  V. 

History — 

One  of  old  metals — called  Stibium  by  Pliny. 

Occurrence — 

Occurs  native  and  in  combination — common  ore  is  stibnite 
(Sb2S3). 

Preparation — 

(1)  Heat  stibnite  with  Fe. 

Sb2S3  + 3Fe  = Sb2  + 3FeS. 

(2)  Heat  stibnite  with  Na2C03  + C. 

2Sb2Ss  + 6Na2C03  + 3C  = 4Sb  + 6Na2S  + 9C02. 

Properties — 

Silver  white  metal,  highly  crystalline,  hard  and  brittle — 
volatizes  at  red  heat — expands  on  cooling — oxidizes  slowly  in 
moist  air,  rapidly  in  HNO3  to  Sb203 — Precipitates  from  its 

(tartaric  acid) 

solution  black  by  Zn — soluble  instrong  HC1  and  H2(C4H40e) 
— insoluble  in  HNO3 — Sb  compounds  are  poisons. 

Used— 

reason  of  its  expansion  on  cooling  Sb  is  used  in  tjrpe 

metal. 

Stibine  (H3Sb)^ 

Analogous  to  H3N  and  H3As — discovered  by  Marsh  process 
— formed  under  same  conditions  as  AsH3,  and  has  same  proper- 
ties— chiefly  important  in  detection  of  Sb. 

ANTIMONY  HALOIDS. 

Sb  forms  two  compounds  with  Cl  and  F — SbCls,  SbCls— 
SI3F3,  SbF5— with  Br  and  I it  forms  SbBr3  and  SM3. 

Antimony  trichloride  (SbCls)— 

A soft  colorless  body— commercially  called  “Butter  of  Anti- 
mony”—prepared: 

(1)  By  heating  in  Cl  gas  (note  PCI3). 

2Sb  + 3C12  = 2SbCl3. 

•(2)  By  dissolving  Sb  in  HC1. 

2Sb  + 6HC1  = 2SbCl3  + 3H2. 


33 


SbCls  is  decomposed  by  H2O  with  precipitation  of  the  oxy- 
chloride or  basic  chloride  SbOCl. 

SbCl3  + H20  = SbOCl  + 2HC1. 

This  white  precipitate  is  a characteristic  reaction  for  anti- 
mony—SbOCl  is  called  “Powder  of  Algaroth.” 

Antimony  pentachloride  (SbCls) — 

Prepared  by  action  of  chlorine  in  excess  on  SbCl3,  and  puri- 
fied by  distilling  in  current  of  Cl — SbCls  is  a fuming  colorless 
liquid,  decomposed  by  water  to  Sb02Cl  Other  haloid  com- 
pounds of  Sb  are  analogous  to  those  of  P and  As. 

OXIDES  OF  Sb. 


Three  compounds— Sb203,  Sb204,  Sb205. 

Sb203  acts  as  base  with  strong  acids. 

Sb204  base  with  strong  acid — acid  with  strong  base. 

Sb205  acts  as  acid. 

Antimony  trioxide  (Sb203) — 

Found  free  in  prismatic  crystals. 

Prepared: 

(1)  Heating  Sb  in  the  air — Sb203  sublimes  as  white  cr}rstals. 

(2)  Oxidizing  Sb  with  HNO3. 

(3)  “ “ “ KN03. 

Sb203  is  isomorphous  with  AS2O3 — acts  as  a base — soluble 
in  strong  HC1,  forming  SbCl3,  which  with  H2O  gives  SbOCl. — 
Antimonic  anhydride  Sb205  is  a pale  yellow  powder  which  re- 
duces to  Sb203  with  heat. 


Antimonic  Acid  H3Sb04— -f 

Analogous  to  H3ASO4.  and  H3PO4, — derived  from  Sb205. 
The  Na  salt  is  insoluble  in  H20  and  of  importance  in  separating 
As  from  Sb. 


Antimony  Sulphides — 

Two  known  sulphides,  Sb2S3  and  Sb2Ss,  analogous  to  sul- 
phides of  As.  The  trisulphide,,  stibnite,  Sb2S3  is  most  important 
ore  of  antimony — separated  from  impurities  by  melting — crude 
Sb2S3  called  “antimony”  oxidizes  with  heattoSb203and  Sb204. 
Antimony  pentasulphite  Sb2Ss  is  an  orange  colored  powder 
formed  by  passing  H2S  into  antimonic  acid. 


34 


BISMUTH— Bi. 

At.  Wt.  210.  Val.  I.  III.  or  V. 

Occurrence — 

Mostly  as  free  Bi  in  Saxony  also  as  Bi203  and  B^S^ — com- 
paratively rare  substance. 

Properties — 

Brittle,  hard,  faint  red- white  color,  with  metallic  lustre — not 
easily  oxidized  but  at  red  heat  con  verted  to  BLO3 — not  attacked 
by  dilute  H2SO4  or  HC1,  but  easily  dissolved  by  HNO3— melting 
point  lowered  by  alloying  with  other  substances,  forming  fusible 
metals,  as  Woods  metal  which  melts  in  hot  water — Bi  does 
not  combine  with  H — unites  with  Cl  and  Bi,  in  two  proportions 
— with  I in  one. 

BISMUTH  CHLORIDES. 

Bi  forms  two  chlorides,  BLCI4  and  BiCL. 

Bismuthous  Chloride  (BLCli)  is  a black,  water-absorbing 
powder,  formed  by  heating  Bi  and  Hg2Cl2- 

2Hg2Cl2  + 2Bi  = 2Hg2  + Bi2Cl4. 

Bismuth  Chloride  (Bids)  is  volatile  like  SbCls  and  formed 
in  same  way,  also  by  dissolving  the  oxide  in  HC1— a soft  unsta- 
ble mass,  decomposed  by  H2O  to  BiOCl. 

Bismuth  oxide  (B^Os) — 

Formed  by 

(1)  Heating  Bi  in  the  air. 

(2)  Heating  Bi0N03. 

A yellow  powder — forms  salts  with  acids  which  are  color- 
less when  the  acid  is  colorless — All  salts  decomposed  by  much 
Hb20 — chief  compound  is  the  basic  nitrate  BiONOs , bismuth 
sub-nitrate,  or  bismuthyl  nitrate,  which  is  much  used  in  medi- 
cine, and  as  cosmetic.  Often  contains  a dangerous  amount  of 
As. 

BORON— B. 

At.  Wt.  11.  Val.  III. 

History — 

Sodium  salt  long  known  as  “Tinkal.” 

Occurrence — 

Not  free— combined  as  boracic  acid  H3BO3,  and  borax  Na2 
B4O7. 


Preparation — 

Treat  Boric  anhydride  with  Na. 

B2O3  + 6Na  = 3Na20  + B2  (amorphous). 

Properties — 

Two  varieties,  amorphous  and  crystalline. 

Amorphous  is  a brown  powder — odorless,  tasteless,  insoluble 
in  H2O,  non-conductor  of  electricity — on  heating  combines  di- 
rectly with  many  substances  asS,  Cl,  Br — Boron  burns  in  air  to 
B2O3  and  N. 

Crystalline  variety  obtained  by  igniting  B with  Cl — thus 
formed  is  hard  as  the  diamond. 

BORACIC  ACID— H3BO3. 

Occurrence — 

In  volcanic  districts,  of  Tuscany  certain  hot  springs  rise, 
called  “fumaroles” — these  carry  H3BO3  mechanically.  Steam 
from  the  fumaroles  is  passed  into  water  and  1%  solution  of 
H3BO3  obtained— this  is  evaporated  by  heat  of  the  springs  and 
purified  through  its  sodium  salt. 

Properties — 

Acid  and  its  salts  form,  on  heating,  a colorless  glass  bead, 
which  dissolves  oxides  with  characteristic  colors. 

H3BO3  itself  forms  no  salts  but  on  heating  it  loses  water, 
condensing  to  other  acids  which  form  salts. 

At  100°  H3BO3  = H20  + HB03. 

Higher  4H3BO3  ==  5H2O  + H2B4O7.  (Borax  = Na2B4.07.) 

(See  Na) 

Still  higher  H2B4O7  = H20  + B203. 

Borates  are  identified  by  a green  color  to  the  flame.  Alkali 
borates  only  are  soluble  in  water. 

Boron  trioxide  (B2O3) — 

The  only  oxide  of  Boron — called  Boracic  anhydride.  3H2O 
+ B2O3  =2H3BC>3.  B2O3  is  the  source  of  all  B compounds,  and 
is  prepared  by  heating  B in  the  air  or  heating  H3BO3. 

2H3BO3  = B2O3  + 3H20. 

Boron  Haloids — 

B unites  with  all  but  I — forms  the  tri-compounds  BiCL.etc. 
— Liquids  of  low  b.p.  and  easily  decomposed  by  H2O— formed 
by  passing  the  haloid  over  heated  B or  B2O3+C. 


SILICON— Si. 

At.  Wt.  28.  Val.  II  or  IV. 


History— 

Isolated  by  Berzelius  in  1823.  Amorphous  variety  found 
first. 

Occurrence — 

Always  combined  with  oxygen  in  Silicon  dioxide  (Si02)  as 
quartz,  sand,  etc.— All  geological  formations  except  chalk  con- 
tain Si. 

Preparation — 

Three  varieties  of  Silicon. 

(1)  Amorphous,  made  by 

(a)  fusing  potassium-fluo-silicate  K2SiF6  with  K or 

NA. 

K2SiF6  + 4K  = Si  + 6KF. 

(b)  use  Silicon  tetra-chloride  and  K. 

SiCl4  + 4K  — Si  + 4KC1  (Berzelius). 

(2)  Graphitic,  by  fusing  the  amorphous  with  Al. 

(3)  Crystalline,  by  fusing  the  graphitic. 

Properties — 

Si  has  strong  affinity  for  O — the  amorphous  variety  burns 
to  SiC>2 — Si  is  soluble  in  Aluminum  (Al)  just  asC  is  in  iron.  Also 
soluble  in  HF  forming  SiF4  + H4. 

Silicic  Hydride  (SiFG) — 

A spontaneously  lighting  gas — burns  to  SiCL  and  HoO— 
analogous  to  methane  (CH4). 

Silicon  Fluoride  (SiF4). 

Prepared  from  CaF2,  SiC>2  and  H2SO4. 

2CaF2  + S1O2  + 2H2SO4  = 2CaS04  + SiF4  + 2HsO. 

Etching  of  glass  depends  on  this  reaction  as  the  Si02  may 
be  in  form  of  glass. 

Si02  + 4HF  = SiF4  + 2H20. 

Properties — 

Colorless  gas — bad  odor— deadly  poison — fumes  with  H2O, 
forming  silicic  acid  and  hydro-fluo-silicic  acid. 

3SiF4  + 3H20  = H2Si03  + 2H2SiF6. 


Hydro=Fluo=Silicic  Acid  (HoSiFo)  — 

Prepared  by  passing  Sip4  into  H2O. 

K salt  is  insoluble  in  HoO. 

Chemical  deportment  similar  to  hydrogen-halogen  acids 
(HC1,  etc). 

Silicon  Dioxide  Si02 — “Silica” — 

At  first  supposed  to  be  an  earth — called  vitreous  earth  from 
use  in  glassmaking — finally  shown  to  be  a weak  acid. 

Found  both  free  and  combined  in  nature.  Free  as  quartz 
(anhydrous),  opal  (hydrous),  amethyst  quartz  (colored  with 
Mn),  agate,  flint,  etc.  In  combination  as  silicates  forming 
most  of  the  “rocks”. 

By  evaporating  to  dryness  with  acids,  silicates  decompose 
with  a residue  oFSiOo  insoluble  in  acids ; this  constitutes  a test 
* for  silica.  Silicates  are  insoluble  save  the  Na  or  K salt. 

Silicic  Acid  H2Si03— 

There  are  many  silicic  acids  but  this  is  the  most  common 
form — Heated  it  loses  water  to  form  SiC>2. 

CARBON— C. 

At.  Wt.  12.'  Val.  II.  or  IV. 

History — 

First  referred  to  as  an  element  by  Lavoisier  1780.  He 
showed  that  “mephitic  air”  or  “carbonic  acid”  contained  C 
and  O,  also  that  C was  the  important  part  of  charcoal. 

Occurrence — 

Free  in  three  modifications — crystalline  as  diamond  and 
graphite — amorphous  as  various  kinds  of  coal — as  CO2  in  air 
and  water — as  carbonates  in  all  soil,  sometimes  forming  moun- 
tains. C crystallizes  easily  as  graphite,  but  not  easily  as  dia- 
mond. Obtained  as  amorphous  C from  organic  bodies. 

Diamond — 

♦ 

The  diamond  is  found  mostly  in  old  rock  strata  chiefly  in 
India,  Borneo,  Brazil  and  South  Africa.  Usually  colorless — 
when  crude  covered  with  dull  crust — when  polished  exhioits 
great  lustre— high  refraction — extreme  hardness — brittle— poor 
conductor  of  heat  and  electricity — not  easily  attacked  by  oxi- 


38 


dizing  agents — unchanged  by  heat  alone — burns  easily  in  oxy- 
hydrogen  blowpipe. 

Carbonado  or  black  diamonds  occur  in  large  masses — have 
hardness  but  not  fire  or  water. 

Graphite — 

Widely  distributed  as  a mineral— called  also  plumbago  and 
“black  lead” — Till  1798  thought  to  contain  Pb — usually  con- 
tains about  five  % ash  (Si02  + Fe203). 

A black  substance — metallic  lustre — very  friable  but  exceed- 
ingly hard — Good  conductor  of  heat  and  electricity — Crystal- 
lizes in  hexagonal  crystals  from  melted  Fe — not  affected  by  O 
at  ordinary  temperatures — at  high  temperature  oxidizes  to 
C02 — insoluble  in  all  ordinary  liquids. 

Used  for  lead  pencils — lubricants — electroU'ping — as  preser- 
vative from  rust — mixed  with  clay  is  used  for  crucibles. 

CHARCOAL. 

Preparation — 

Not  easily  obtained  pure — made  by  distilling  wood  in  re- 
torts or  by  burning  wood  in  small  amount  of  air,  getting  rid 
of  volatile  substances  and  saving  most  of  the  carbon. 

Properties — 

Charcoal  is  brittle  because  of  cell  structure  of  wood  which 
is  retained. 

Physical  properties  are  important — absorbs  and  condenses 
great  varietj^  of  gases  and  vapors,  especially  such  as  dissolve 
in  H20 — action  is  not  mere  absorption,  but  oxidation  of  offen- 
sive gases,  hence  charcoal  is  a valuable  disinfectant,  by  hasten- 
ing destruction  of  putrescible  organic  matter — is  not  an  anti- 
septic or  preservative  agent — C likewise  decolorizes,  being  es- 
pecially used  in  purification  of  sugar,  and  in  chemical  and  phar- 
maceutical preparations. 

Bone  Black — 

Prepared  by  destructive  distillation  of  bones — particularly 
good  for  all  uses  of  charcoal  because  of  large  surface. 

Coke— 

By  destructive  distillation  of  soft  coal — chemically  classed 
between  graphite  and  charcoal. 


— 39- 


Lam  p=black — 

The  kind  of  charcoal,  from  burning  oils  in  small  amount  of 

air. 

Gas  Carbon- 

Carbon  of  gas  retorts  is  formed  by  long  heating  of  the  crust 
which  condenses  on  the  interior  of  gas  retorts — very  hard,  com- 
pact and  dense— conducts  heat  and  electricity — employed  in 
galvanic  batteries  and  as  pencils  for  electric  lamps.  Chemically 
classed  between  charcoal  and  graphite. 

Coal- 

Chief  varieties  are  Anthracite,  Bituminous  and  Peat. 

The  amount  of  carbon  decreases  while  water  and  S increase 
in  the  given  order. 

CARBON  AND  HYDROGEN. 

Form  an  immense  number  of  compounds  known  as  hydro- 
carbons. One  is  formed  by  direct  union  of  elements — C2H2, 
acetylene — formed  with  C electrodes  in  atmosphere  of  H. 

METHANE— CH4  (MARSH  GAS). 

Occurrence — 

Found  where  vegetable  matter  decays  under  water — hence 
called  Marsh  gas — Found  also  in  coal  mines  where  it  often 
causes  explosions,  hence  named  “fire  damp” — chief  constituent 
of  gas  wells. 

Preparation — 

Distilling  NaC2Hs02  with  NaOH — 

NaC2H302  + NaOH  - Na2C03  + CH4. 

Synthetically  by  CS2  + H2S  + red  hot  Cu — 

CS2  + 2 H2S  + 8 Cu  = CH4  + 4 Cu2S. 

Properties — 

Colorless,  inodorous,  tasteless  gas  (condensed  by  188  at.) 
— slightly  soluble  in  water — burns  with  slight  luminous  flame — 
mixture  with  air  explodes  (“fire  damp”),  forming  H2O  + CO2, 
“choke  damp.” 

CARBON  MONOXIDE— CO. 

Properties — 

A colorless  gas,  inodorous,  insoluble  in  H2O — burns  with 


— 40- 


pale  blue  flame  to  CO2 — highly  poisonous— a powerful  reducing 
agent,  forming  CO2. 

Preparation— 

Always  formed  by  incomplete  combustion  of  C.  At  first 
prepared  by  distilling  ZnO  with  C and  supposed  to  contain  H. 

a)  Pass  steam  over  red-hot  coal. 

c + H20  = CO  + h2. 

This  is  “water  gas,”  one  form  of  illuminating  gas — the  other  or 
coal  gas,  formed  by  destructive  distillation  of  coal,  is  less  poi- 
sonous. 

b)  CO  is  formed  by  decomposition  of  many  organic  sub- 
stances, especially  oxalic  acid. 

H2C2O4  = C02  + CO  + H20. 

CARBON  DIOXIDE— C02. 

History — 

Known  in  earliest  times  through  action  of  vinegar  on  chalk 
—first  distinguished  as  a gas  by  V.  Helmont,  who  called  it  “gas 
sylvestre” — Lavoiser  termed  it  carbonic  acid  and  showed  its 
composition. 

Occurrence — 

Widely  distributed  in  atmosphere  and  soil — occurs  in  liquid 
form  in  some  crystalline  substances — generally  found  united 
with  bases,  especially  in  CaC03. 

Preparation — 

By  action  of  acids  on  carbonates — 

CaC03  + 2HC1  = CaCl2  + C02  + H20. 

Properties — 

Colorless  gas — acid  odor — heavy — incombustible — extin- 
guishes flame — will  not  support  respiration — liquid  form  is  col- 
orless— becomes  solid  on  rapid  evaporation,  and  forms  with 
ether  a freezing  mixture — CO2  is  absorbed  most  by  water  at 
low  temperature,  making  carbonic  acid  waters — its  presence 
in  wine  and  beer  is  due  to  fermentation — true  carbonic  acid  is 
H2CO3,  a very  unstable  substance  forming  two  classes  of  salts — 
carbonates  of  alkalies  only  are  stable,  but  acid  salts  of  alkalies 
are  decomposed  by  heat. 

2NaHC03  = Na2C03  + H20  + C02. 


CARBON  DISULPHIDE— CS2. 


History — 

Discovered  1796  by  Lampadius  and  called  “sulphur  alco- 
hol.” Thought  to  contain  S,  H,  C,  N,  but  in  1811  shown  to  be 
only  C and  S. 

Preparation — 

Pass  vapor  of  S over  red-hot  C. 

Properties — 

Heavy,  colorless,  strongly  refracting  liquid — etherial  odor 
when  fresh — volatile,  with  poisonous  vapor,  having  cumulative 
effect — easily  ignited  and  burns  to  C02  + S02 — vapor  bums 
with  NO — mixed  with  air  its  vapor  explodes — remarkable  anti- 
septic power — Decomposed  by  heat  to  its  elements — CS2  forms 
salts  of  thiocarbonic  acid  H2CS3. 

CS2  used  chiefly  as  solvent  especially  of  S and  P,  also  of  Br 
and  I. 

CYANOGEN— C2N2. 

H istory — 

Discovered  in  1815  by  Gay-Lussac  while  investigating 
“prussic  acid” — called  it  cyanogen  from  prussian  blue. 

Preparation — 

C and  N unite  only  with  difficulty — the  compounds  are  ob- 
tained indirectly  and  all  contain  the  group  CN,  which  was  the 
first  “radicle”  discovered. 

1)  C2N2  formed  when  organic  bodies  are  heated  with  K or 
Na — detection  of  N in  organic  bodies  depends  on  this. 

2)  Heat  Hg(CN)2  = Hg  + C2N2. 

Properties — 

Colorless  gas  — strong  odor — intensely  poisonous — burns 
with  red  flame  to  C02  + N ; CN  is  recognized  by  its  formation 
of  Prussian  blue — when  in  free  state  the  radicle  (CN)  exists  as 
C2N2  Dicyanogen,  which  closely  resembles  the  Haloids. 

HYDROCYANIC  ACID-HCN 

Occurs  in  nature  in  kernels  of  almonds,  peaches,  plums,  etc.; 
blossoms  and  leaves  of  the  peach  tree,  and  several  other  trees 


•4.2 


and  shrubs— a liquid  similar  in  properties  to  HC1,  but  intensely 
poisonous — generally  prepared  by  treating  K salt  with  HC1, 
KCN  + HC1  = HCN  + KC1. 

ACETIC  ACID— H(C2H302) 

Pyroligneous  acid  a bi-product  in  making  charcoal — clear, 
colorless  liquid — characteristic  taste  and  odor.  Pure  acid  is 
called  “glacial” — dilute,  impure  form  is  known  as  vinegar — 
H(C2H302)  is  a monobasic  acid  and  forms  acetates  M(C2H302). 

OXALIC  ACID— H2C204. 

Crystalline  white  solid — poisonous — good  reducing  agent. 
H2C204  is  a dibasic  acid  formed  by  action  of  HN03  on  many 
organic  substances,  especially  the  sugars,  starch,  etc. — whole- 
sale by  action  of  KOH  on  sawdust — occurs  widely  distributed, 
especially  in  plants  of  oxalis  variety  and  in  urinary  calculi. 

CHROMIUM— Cr. 

At.  Wt.  52.4.  Yal.  II  or  III. 

History — 

In  1797  was  discovered  by  Vanquelin  in  mineral  “crocoisite” 
(PbCr04). 

Occurrence — 

Not  widely  distributed — chief  ore  is  “chromite”  or  chrome 
iron  ore  (FeO  * Cr203). 

Extraction — 

Ore  is  pulverized  and  heated  with  K2C03  and  CaO — K2Cr04 
potassium  chromate  is  formed  and  extracted  with  water — 
H2S04  is  added  to  form  K2Cr207  potassium  bichromate, 
which  is  the  commercial  salt.  The  metallic  Cr  is  obtained 
by  heating  Cr203  (chromic  oxide)  with  C. 

Properties — 

A hard,  difficultly  fusible  metal — similar  to  iron — with  heat 
it  slowly  oxidizes  to  Cr203 — burns  in  oxygen  with  bright  light 
— soluble  in  HC1  and  H2S04 — not  altered  by  HN03— forms 
chromous  and  chromic  compounds — Cr  forms  an  immense  num- 
ber of  compounds,  in  some  it  is  acidic,  in  some  basic,  or  its  acid 
and  basic  oxides  may  unite,  forming  a salt  of  itself. 


43 


CHROMIUM  HALOIDS— CrCl2,  Cr2Cl6,  CrF6. 

Chromous  chloride — 

A white  powder  soluble  in  water — formed  by  reducing 
Cr2Cl6  with  H — chromous  compounds  are  hard  to  obtain  and 
easily  oxidize  to  chromic. 

Chromic  Chloride  Cr2Cle — 

A violet  sublimate  obtained  by  ignition  of  Cr203  with  C 
•in  a current  of  Cl — insoluble  when  pure  but  easily  dissolves  if  a 
trace  of  CrCl2  is  present — the  hydrated  salt  Cr2Cl212H20 
forms  green  deliquescent  crystals. 

Chromic  Fluoride  CrFe — 

Chiefly  interesting  as  showing  the  sexivalent  character  of  Cr. 
CHROMIUM  OXIDES. 

CrO — Cr203 — Cr03 — first  two  basic — last  acidic. 

Chromous  Oxide  CrO — 

Known  only  in  its  Hydrate  Cr(OH)2,  which  is  made  from 
CrCl2  + KOH  and  rapidly  changes  to  chromic  oxide. 

Chromic  Oxide  Cr203 — 

Found  in  nature  as  “chromite” — formed  as  a green  powder 
by  ignition  of  ammonium  di-chromate 

(H4N)2Cr207  = Cr20  + 4H20  + N. 

The  ignited  oxide  is  insoluble  in  acids — fused  with  silicates  it 
gives  them  an  emerald  green  color — used  to  color  glass  and 
porcelain — when  freshly  precipitated  the  hydrate  is  soluble  in 
acids.  It  is  used  as  a paint,  “Gingnet’s  green” — like  all  sesqui- 
oxides,  Cr203  is  normally  basic,  but  will  not  give  salts  with 
weak  acids.  On  the  contrarv,  with  strong  bases  it  exhibits  an 
acid  character,  forming  chromites. 

Chromic  anhydride  Cr03— 

Red  deliquescent  crystals — active  oxidizing  agent — highly 
corrosive — made  from  the  K chromate  or  bichromate  by  treat- 
ing with  H9SO4 — 

K2Cr04  + H2S04  = Cr03  + H20  + K2S04 
The  water  solution  contains  the  unstable  chromic  acid,  H2Cr04, 
which  is  analogous  to  H2S04  but  known  only  through  its  salts 
the  chromates  and  bichromates — the  neutral  chromates  are  gen- 


44 


erally  3^ellow,  the  bichromates  red — basic  Cr203  unites  with 
acidic  CrOa  to  form  chromium  chromate,  Cr2Cr04 — The  Ba, 
Pb,  A g,  and  Hg  salts  of  chromic  acid  are  insoluble. 

Chrome  alum  Cr2'K2(S04)4  + 24H20— 

The  chief  salt  where  Cr  acts  as  a base — formed  as  violet 
octohedra,  when  a solution  of  K2Cr207  -f  H2S04  is  treated 
with  S02,  and  evaporated  below  80° — above  80°  the  solution 
turns  green  and  refuses  to  crystallize. 

Chromium  sulphides  CrS  and  Cr2S3 — 

Cr  forms  two  sulphides,  CrS  and  Cr2S3  corresponding  to 
the  oxides  but  cannot  be  formed  in  the  wet  way. 

MOLYBDENUM-MO. 

At.  Wt.  96.  Val.  II,  IV,  or  VI. 

H istory — 

Name  comes  from  the  Greek  for  graphite,  with  which  the 
mineral  Molybdenite  was  confused — M0O3  was  separated  in 
1778  by  Scheele  and  Mo  in  1798  by  Hjelm. 

Occurrence — 

Chief  ore  is  molybdenite  MoS2. 

Properties — 

Hard,  silver  white  metal — very  stable — when  pure,  infusible 
— after  long  heating  converted  to  Mo203 — dissolves  in  HNO3 
and  aqua  regia. 


COMPOUNDS  OF  Mo. 

Chief  compounds  are  the  acid  H2Mo04  and  its  (NH4)  salt 
(NH4)2Mo04. 

Jlolybdic  Oxide  M0O3 — 

The  anhydride  of  mofybdic  acid  is  M0O3,  made  by  roasting 
molybdenite — it  combines  readily  with  bases,  forming  molyb- 
dates. 

COMPOUNDS  OF  Mo. 

Molybdic  acid  H2Mo04 — 

Is  formed  wnen  molybdates  are  treated  with  dilute HNO3 — 
Dissolved  in  strong  ammonia  it  forms  (H4N)2M04,  ammonium 
molybdate,  which  in  HNO3  solution  is  used  to  precipitate 


H3PO4  in  the  form  of  ammonio-phospho-molybdate,  a delicate 

TEST  for  PHOSPHATES. 


TUNGSTEN— W. 

At.  Wt.  184.  Val.  II,  IV,  or  VI. 

Analogous  to  Mo — forms  similar  compounds — Mo  com- 
pounds crystallize  easier — Tungstic  acid  similar  to  molvbdic 
acid — common  salts  are  Na  and  H4N  tungstates — the  former 
used  to  make  fabrics  fire-proof — chief  ore  is  wolframite,  FeWC>4. 
Calcium  tungstate,  CaWC>4,  is  used  as  a fluorescent  screen. 

TIN-Sn. 

At.  Wt.  118.  Val.  II  or  IV. 

H istory — 

Used  by  the  Phoenicians,  who  got  it  from  England. 

Occurrence — 

Native  in  small  amounts — chief  ore  is  tin  stone  SnOo. 

Extraction — 

By  reducing  the  ore  with  carbon  and  remelting. 

Properties — 

Lustrous  white  metal — crystalline — when  bent  emits  a pe- 
culiar sound  (tin  cry) — ductile — malleable,  but  more  so  at  100° 
— brittle  at  200° — not  oxidized  by  the  air  till  its  melting  point 
— at  white  heat  burns  brilliantly  to  SnC>2 — forms  salts  with  HC1 
and  H2SO4 — forms  SnCU  with  HNO3  or  with  the  alkalies,  when 
it  acts  as  an  acid. 

sodium  stannate 

2 Sn  +-6NaOH  = 2Na3Sn03  + 3H20. 

Sn  is  not  affected  by  H2S. 

Use — 

Chiefly  in  “tinware”  which  is  iron  covered  with  a layer  of 
tin — also  in  bronze  (Cu,  Sn,  Zn),  solder  (Sn,  Pb),  Brittania 
metal  (Sn,  Sb),  tin  amalgam  (Sn,  Hg) — the  last  is  used  in  sil- 
vering mirrors. 

TIN  CHLORIDES. 

Stannous  chloride  SnCL — 

The  anhydrous  form  is  made  by  HC1  gas  on  metallic  tin — 
hydrated  by  treating  excess  of  Sn  with  HC1 — SnCU  comes  on  the 


46 


market  as  tin  salt — is  a white  crystalline  substance — soluble 
in  small  amount  of  H2O,  in  large  amount  forms  the  basic  chlor- 
ide— has  a strong  tendency  to  unite  with  Cl,  hence  used  as  a re- 
ducing agent  to  form  SnCU- 

Stannic  Chloride  SnCU — 

A colorless,  fuming  liquid,  hence  name  of  “Spiritus  fumans 
Libavii” — made  by  action  of  Cl  on  Sn  or  SnCU — combines  with 
metallic  chlorides  to  form  double  salts,  of  which  “pink  salt” 
SnCU : 2H4NCI  is  most  important — both  chlorides  of  tin  are 
used  as  mordants — the  bromides  and  iodides  are  analogous. 

TIN  OXIDES. 

Stannous  Oxide  SnO — 

The  basis  of  stannous  salts — formed  by  heating  the  hydrate 
— SnO  is  a white  powder  which  easily  oxides  to  Sn02. 

Stannic  Oxide  Sn02 — 

“Tinstone”  is  a crystalline  solid  formed  by  heating  Sn  or 

SnO — insoluble  in  acids fused  with  NaOH  or  KOH  it  forms 

soluble  stannates — Sn02  acts  also  as  a weak  base  forming  with 
H2SO4  stannic  sulphate  Sn  (SO4U 

ACIDS  OF  TIN. 

Stannic  Acid  H2Sn03— 

Separates  as  a white  precipitate  when  HC1  is  added  to  po- 
tassium stannate  (K^SnOs) — dissolves  easily  in  HN03,HC1  and 
the  alkalies — left  under  water  or  in  vacuo,  stannic  gradually 
changes  to  metastannic  acid,  an  insoluble  modification  of  the 
same  composition.  This  isomer  is  also  formed  when  tin  is 
treated  with  HNO3,  concentrated — metastannates  are  entirely 
different  from  stannates. 

Stannic  Phosphate  Sn3(P04)2 — 

Tin  forms  with  phosphoric  acid  a compound,  Sn3(PCU)2> 
which  is  insoluble  in  HNO3 — used  in  analysis  to  remove  H3PO4. 

TIN  SULPHIDES. 

Stannous  Sulphide  SnS — 

Formed  by  direct  union  at  high  temperature — also  precipi- 
tated from  stannous  salts  by  H2S  in  brown  amorphous  form. 


47 


Stannic  Sulphide  S11S2 — 

Precipitated  as  a yellow  powder,  by  H2S,  from  stannic 
salts — when  sublimed  it  forms  a bright  yellow  crystalline  mass 
known  as  mosaic  gold  which  is  used  in  bronzing. 

STANNOUS  SALTS. 

If  acid  is  colorless  the  salt  is  either  colorless  or  yellow — 
stannous  salts  have  metallic  taste — absorb  O from  the  atmos- 
phere— change  easily  to  stannic — 

KOH  gives  white  Sn(OH)2  soluble  in  excess. 

H4N(0H)  gives  white  Sn(OH)2  insoluble  in  excess. 

AuCL  gives  purple  of  Cassius,  characteristic. 

HgCL  gives  Hg2Cl2,  characteristic. 

TITANIUM  Ti— ZIRCONIUM  Zr— THORIUM  Th. 

At.  Wt.  48.  Yal.  IV.  At.  Wt.  91.  Val.  IV.  At.  Wt.  232.  Val.  IV. 

Three  rare  metals  closely  resembling  tin  but  more  basic  in 
character  and  forming  NCkous  salts — The  most  important  is  Ti 
which  generally  accompanies  iron,  especially  in  Titanic  Iron, 
FeTi03 — metal  is  obtained  by  decomposing  potassium  flno- 
titanate,  K^TiFe,  with  K (see  silicon) — a magnetic  dark  gray 
powder — burns  in  air  and  chlorine. 

Ziconium  closely  follows  Ti  but  has  been  obtained  in  crys- 
talline form,  which  looks  like  antimony  but  is  harder — the  oxide 
Zr02  is  used  for  lime  light. 

POTASSIUM— K. 

At.  Wt.  39.— Val.  I 

History— 

“Potash”  from  ashes  known  to  ancients — not  distinguished 
from  “Soda”  till  middle  of  18  century— metal  K isolated  by 
Davy  in  1807. 

Occurrence — 

The  salts  are  found  in  rocks  and  all  cultivated  ground — In 
plants  K occurs  as  oxalate  and  tartrate — In  animals  as  chlor- 
ide and  phosphate — Sweat  on  sheeps  wool  is  one  source  of  K 
compounds — Largely  obtained  from  “Argol”  an  acid,  potas- 
sium tartrate  (HKC4H4O6)  which  mixed  with  coloring  matter 


48 


deposits  on  side  of  wine  casks — Another  source  is  plant  ashes 
which  yield  a crude  K carbonate  called  potash — At  Stassfort, 
large  amounts  are  obtained  from  mineral  carnallite  (KC1). 

Preparation — 

When  Argol  is  heated  in  a closed  retort  it  breaks  up  to  cal- 
cium and  potassium  carbonate  and  carbon — on  farther  heating 
metallic  K distills  as  a green  vapor — essential  reaction  is 
K2C03  + 2C  = 2K-f  3C0 

Properties — 

A soft  bluish  white  metal — green  vapor — phosphoresces  in 
dark — Of  all  the  elements  K has  the  greatest  affinity  for  oxygen 
and  chlorine  at  ordinary  temperatures — it  burns  with  a violet 
blue  flame,  but  oxidizes  in  the  air  without  taking  fire — oxidizes 
on  water,  and  decomposes  it. 

H20  + K = K20  + Ho. 

K20  + H20  = 2K0H. 

K and  Sodium  are  used  to  reduce  metals  from  their  oxides. 

Potassium  Oxide  K2O — 

The  only  important  oxide  of  Potassium  is  K^O,^  gray  solid 
formed  by  action  of  K on  KOH — difficult  to  secure  as  it  eagerly 
unites  with  HoO  forming  KOH. 

Potassium  Hydrate  KOH — 

Formed  when  K acts  on  H2O. 

Generally  prepared  by  treating  a salt  of  K with  hydrate  of 
some  metal  which  will  form  an  insoluble  salt  with  the  radicle 
of  the  K salt. 

K2C03  + Ca(OH)2  = 2K0H  + CaC03— insoluble. 

Crude  KOH  obtained  from  wood  ashes. 

KOH  is  best  example  of  an  alkali  and  is  strongest  of  all 
bases — it  neutralizes  acids  to  form  salts — decomposes  fats  and 
oils  to  soaps  hence  destroys  animal  tissue— is  soluble  in  V2  its 
weight  of  water,  forming  a crystalline  hydrate  KOH  * 2H2O. 

KN03  (Saltpeter). 

Occurrence — 

Spoken  of  under  HNO3. 


4-9 


Preparation — 

Usually  by  treating  Chili  saltpeter  NaNOs  with  KCL 

NaN03  + KC1  = KNOs  + NaC.l. 

NaCl  being  less  soluble  is  removed  by  evaporation. 

Properties — 

Crystallizes  in  white,  anhydrous,  striated  prisms,  with  cool- 
ing taste,  and  produces  cold  by  its  solution — large  amounts  act 
as  poison — heated  yields  KNO2  and  oxygen,  heated  higher  de- 
composes to  KOH. 

2KN03  = K20  + 2N02  + 0. 

I<2  + H20  = 2K0H. 

Melted  in  drops  is  “sal  prunelle.” 

KNO3  oxidizes  most  substances  save  a few  metals — with  C 
it  deflagrates. 

4KNO3  + 5C  = 2K2C03  + 3C02  + 2N2. 

Use — 

Used  in  medicine  and  preparation  of  HNO3. 

Chief  use  is  in  making  gunpowder.  Used  instead  of  NaNOa 
because  KNO3  is  not  deliquescent.  The  theoretical  equation  in 
discharge  of  gunpowder  is — 

2KN03  + S + C3  te  K2S  + N2  + 3C02. 

Theoretical  proportions  are — 

, KNOs  s c 

74.8  11.8  13.4  parts. 

POTASSIUM  HALOIDS. 

2 Potassium  chloride  KC1 — 

Called  digestive  salt,  also  sal  febrifugium  sylvii — found  in 
sea- water  and  springs — occurs  as  sylvite  and  earn  alii  te  at 
Stassfurt — prepared  from  K2CO3  and  HC1— 

K2C03  + 2 HC1  = 2 KC1  + C02  + H20. 

Crystallizes  in  colorless  cubes — easily  soluble  in  H2O — solution 
absorbs  heat — used  largety  in  making  alum,  and  K2CO3 — crude 
salt  used  as  a fertilizer. 

Potassium  bromide  KBr — 

Colorless  cubes — soluble  in  H2O — made  (1)  directly  from 
elements.  (2)  By  action  of  Br  on  KOH — 

6 KOH  + 6 Br  — 5 KBr  + KBr03  + 3 H20. 

KB1O3  = KBr  + 0. 

(Ferrous  Bromide)  (Ferrous  Carbonate) 

Usually  (3)  FcBr2  + K2C03  = 2 KBr  + FeC03. 

Used  in  medicine. 


0 


Potassium  Iodide  KI — 

White  crystalline  solid,  used  in  medicine  and  photography 
—usually  made  by  action  of  K2CO3  on  HI  or  Fel2~ 

(1)  K2CO3  + 2HI  = 2KI  + H20  + C02. 

(2)  K2CO3  + Fel2  - 2 KI  + FeO  + C02. 

Potassium  Chlorate  KCIO3  — 

White  crystalline  solid  made  by  treating  KOH  with  Cl — 

6 KOH  + 6 Cl  - 5 KC1  + KC103  + 3 H20. 

Substance  treated  under  chlorine. 

Potassium  Perchlorate  KCIO4 — 

White  solid  formed  by  decomposition  of  KCIO3 — 

2 KCIO3  = KCIO4  + KCI  + 02. 

Notable  for  its  slight  solubility. 

Potassium  Sulphide  K2S — 

Potassium  and  sulphur  form  many  compounds  all  soluble 
in  water — chief  is  K2S — prepared  by  fusing  K2SO4  with  C. 

K2SO4  + 2C  = K2S  + 2C02. 

When  fused  is  a red  mass,  but  crystallizes  from  water  solu- 
tions in  colorless  deliquescent  prisms — “Hepar  Sulphuris”  or 
“liver  of  sulphur”  is  a mixture  of  the  Polysulphides  of  K (K2S, 
K2S4,K2S5)  and  K2SO4,  obtained  by  fusing  K2CO3  + S. 

Potassium  Sulphydrate  KHS — 

White, crystalline, solid,  with' alkaline  reaction,  prepared  by 
action  of  H2S  on  KOH. 

KOH  + H2S  = KHS  + H20. 

Potassium  Sulphate  K2SO4 — 

White,  crystalline,  solid — native  as  Kainite — prepared  from 
K2CO3  + H2SO4  = K2SO4  + h2o  = co2. 

Soluble  in  H 2 0— insoluble  in  absolute  alcohol.  Kainite  is 
used  as  fertilizer. 

Potassium  Nitrite  KNO2 — 

Formed  b\r  reduction  of  KNO3  by  heat. 

Used  as  source  of  HNO2  and  in  analysis. 

Potassium  Arsenite  K3ASO3 — 

Formed  by  action  of  K2CO3  011  AS2O3. 

A dilute  solution  used  in  medicine  under  name  of  “Fowler’s 
solution.” 


•51 


Potassium  Pyro=Antimoniate  K2HoSb07— 

Made  from  KN03  and  Sb. 

Used  in  testing  for  Na. 

Potassium  Carbonate  K2C03— 

Deliquescent  salt— strongly  alkaline — insoluble  in  alcohol. 
Prepared  by  ignition  of  K salts  of  organic  acids.  Large 
quantities  obtained  from  wood  ashes— also  from  KC1  as  in  the 
LeBlanc  process  for  sodium  carbonate,  which  see. 

Potassium  Hydrogen  Carbonate  KHCO3— 

Called  “Bicarbonate  of  potash”  and  prepared  bypassing 
CO2  through  a solution  of  K2CO3 — 

K2CO3  + C02  + H20  = 2 KHCO3. 

KHCO3  is  less  soluble  than  K2CO3  and  its  solution  o-jves  a 
neutral  reaction. 


Potassium  Cyanide  KCN— 

A poisonous  salt— treated  under  Cyanogen. 
Potassium  Thiocyanate  KCNS— 

Formed  by  treating  molten  KCN  with  S— 
KCN  + S - KCNS. 

Used  as  a reagent  for  Iron  (Fe). 


SODIUM— Na-I. 

History — 

Metal  discovered  in  1807  by  Davy. 

Occurrence — 


Not  free  found  in  many  minerals— occurs  especiallv 
NaCl  m sea  water,  springs  and  as  rock  salt,  and  as  NaN03 
Chih  saltpeter. 

Preparation — 


Metal  prepared  as  K,  but  easier,  hence  cheaper. 

Properties — 

Similar  to  potassium— silver-white  metal  with  purple  vapor 
affinity  for  Cl,  O,  etc.,  only  little  less  than  that  of  K—  decom 
poses  HsO  with  explosion*,  and  if  water  is  warm  burns  with  a 
yellow  flame-forms  with  K the  only  liquid  alloy  not  contain- 
ing Hg— alloy  has  same  appearance  as  Ho- 
-M  j >U<3L  “>• 


of  melted  NaOH  finally 


touching  the 


h2o, 


SODIUM  CARBONATE— Na2C03- 10  H20. 

Occurrence — 

Found  in  nature  in  sea-plants  just  as  K2C03  is  found  in  land 
plants. 

Preparation — 

Formerly  obtained  from  ashes  of  sea  plants — may  be  gotten 
by  heating  a solution  of  the  primary  carbonate,  HNaC03 — 

2 HNaCOs  = Na2C03  + 2 C02  + 2 H20. 

Commercially  in  two  ways — 

(1)  LeBlanc’s  process  from  NaCl. 

(a)  NaCl  to  Na2S04  by  heating  with  H2S04 — 

2 NaCl  + H2S04  = Na2S04  + 2 HC1. 

(b)  Na2S04  to  Na2S  by  heating  with  C— 

Na2S04  + 2 C — Na2S  -f-  2 C02. 

(c)  Na2S  to  Na2C03  by  heating  with  CaC03 — 

Na2S  + CaC03  = Na2C03  + CaS. 

CaS  is  a waste  product. 

2)  Ammonia  or  Solvay  process,  which  is  mor£  modern— 

Mono-am.  carbonate. 

(a)  NaCl  to  HNaC03  by  treating  with  HNH4  C03. 

NaCl  + HNjHf  03  = HNaCOs  + NH4C1. 

(b)  HNaC03  to  Na2  C03  by  heating. 

2HNaC03  = Na2C03  + C02  + H20. 

The  C02  is  passed  into  ammonia  forming  again  the  acid 
ammonium  carbonate  HNH4C03. 

The  NH4CI  in  (a)  is  heated  with  lime  or  magnesia  and  NH3 
set  free  as  H^NOH — into  ammonia  water  thus  formed,  C02from 
(b)  is  run  forming  HNH4C03  again. 

The  Solvay  process  has  no  troublesome  residue  like  theCaS 
of  Le  Blanc. 

Properties — 

Na2C03  is  an  alkaline  salt,  efflorescent  and  very  soluble  in 
H20. 

Use- 

Essential  in  glass  and  soap  making  and  a reagent  in  the 
labor  at  ory. 


53 


Sodium  Sulphate  Na2S04. . IOHoO — 

‘‘Glauber’s  salt”  occurs  in  nature  especially  in  certain  nat- 
ural w aters,  as  Carlsbad  and  Freidrichsbad  springs— is  a by- 
product in  making  HNO3. 

2NaNOs  + H2SO4  = Na2S04  + 2HN03. 

Crystallized  in  large  colorless  monoclinic  crystals  easily  soluble 
in  water  and  easil}^  forming  supersaturated  solutions. 

Used  as  a purgative  in  medicine  and  for  producingartificial 
cold  in  the  laboratory. 

Sodium  Thiosulphate  Na2S2C>3 — 

“Hypo”  is  a colorless  crystalline  solid  easily  soluble  in  H2O 
— may  be  made  by  adding  S to  boiling  solution  of  Na2SC>3. 
Na2S03  + S = Na2S203. 

Chiefly  used  in  photography  and  in  bleaching  as  “anti-chlor.” 

Sodium  Nitrate  NaNC>3 — 

“Chili  saltpeter”  is  a deliquescent  salt  (see  KNO3)  of  value 
as  the  source  of  HNO3  and  KNO3,  also  used  in  coarser  grades 
of  gunpow  der. 

Di=sodium  Phosphate  Na2HP(>4/  I2H2O — 

White  rhombic  prisms — soluble  in  H2O  with  slight  alkaline 
reaction — formed  from  Na2C03  + H3PO4 — 

Na2C03  + H3P04  = Na2HP04  + C02  + H20. 

The  crude  H3PO4  containing  CaSC>4  and  H2SO4  can  be  used 
because  Na2SC>4  will  not  crystallize  together  with  Na2HPC>4. 

Acid  Sodium  Phosphate  NaH2P04‘H20— 

White  rhombic  prisms  easily  soluble  in  H2O,  with  strong 
acid  reaction — chief  cause  of  acid  reaction  in  urine — made  from 
Na3PC>4  and  H3PO4 — 

Na3P04  + 2H3P04  = 3NaH2P04. 

Sodium  Arsenite  Na3As03 — 

An  uncrystallized  substance  formed  from  AS2O3  and  NaOH. 
As203  + 6NaOH  = 2Na3As03  + 3H20. 

Sodium  Arseniate  Na3As04— 

Arseniate  of  soda  forms  in  large  crystals — prepared  by  igni- 
tion of  AS2O3  and  NaOH  with  NaN03. 

Used  in  dissecting  room  as  an  antiseptic. 


Sodium  Tetraborate  NaL>B407 . 10H2O— 

Borax  is  formed  as  “tinkal”  in  India  and  California,  but  is 
mostly  made  artificially  in  Tuscany — heated,  it  loses  water — 
melted,  dissolves  most  of  the  metallic  oxides,  hence  its  use  in 
soldering,  bead  and  blowpipe  work. 

Sodium  Oxides — 

Na  forms  oxides  like  K,  chief  being  Nai>0,  Sodium  oxide, 
which  in  preparation  and  properties  resembles  K2O. 

Sodium  Hydrate  NaOH — 

Closely  resembles  KOH.  Pure  is  fjrepared  from  Na  + PUO, 
commercially  from  Na2COs  + Ca(OH)2. 

Na2C03  + Ca(OH)2  = 2NaOH  + CaC03. 

Purified  by  dissolving  in  alcohol. 

Sodium  Chloride  NaCl — 

“Common  salt”  is  similar  to  KC1 — Crystallizes  in  cubes 
without  water — reaction  of  its  solution  is  neutral — temperature 
makes  but  slight  difference  in  its  solubilitj^ — slightly  soluble  in 
alcohol — insoluble  in  HC1  con. 

Sodium  Hydrogen  Carbonate  NaHCOs — 

Saleratus  or  bicarbonate  of  soda  is  less  soluble  than  Na2C03 
and  is  prepared  by  running  CO2  into  a solution  of  carbonate — 
Na2C03  + H20  + C02  = 2 HNaC03. 

Used  in  medicine  and  preparation  of  effervescing  drinks. 

Sodium  Sulphide  Na2S — 

Most  important  sulphide  is  Na2§,  sodium  sulphide — analo- 
gous to  K2S — has  alkaline  reaction — formed  by  reduction  of 
Na2SC>4  with  C — 

Na2S04  + 4 C = Na2S  + 4 CO. 

Occasionally  used  in  analysis. 

Detection  of  Na  salts — 

Na  salts  detected  by  the  yellow  sodium  flame  or  by  precipi- 
tating as  sodium  antimoniate.  All  Na  salts  except  NaNOs  are 
efflorescent. 


AMMONIUM— H4N. 

History — 

The  group  H4N  has  not  been  isolated,  as  it  breaks  up  to 
H3N  + H — after  the  discovery  ofNa  and  K attempts  were  made 


at  H4N,  but  were  unsuccessful — by  decomposition  of  H4N  salt 
with  Hg  electrode,  ammonium  forms  a voluminous  unstable 
amalgam  with  Hg. 

AMMONIUM  CHLORIDE  H4NCe. 

History — 

“Sal  ammoniac”  was  first  made  in  Egypt  by  sublimation 
of  camel’s  dung. 

Occurrence — 

Small  amounts  found  native  in  volcanic  regions 

Preparation — 

Formed  from  NH3  + HC1  = (H4N)C1. 

Obtained  from  the  liquors  of  gas  works  (see  NH3) — the  gas> 
water  is  treated  with  H2SO4. 

2H3N  + H2S04  = (H4N)2S04. 

( H4N ) 2SO4  is  then  sublimed  with  NaCl. 

(H4N)2S04  + 2NaCl  = 2H4NC1  + Na2S04. 

Properties — 

Colorless,  octohedral,  crystals,  highly  soluble  in  water — caus- 
ing cold  by  solution — volatile  when  heated  by  its  dissolution 
into  NH3  + HC1.  H4NCI  united  easily  with  other  salts  to  form 
double  salts 

The  other  haloid  salts  of  H4N  are  analogous,  to  H4NCI. 
Used  in  medicine  and  the  arts. 

H4N  SALTS. 

Sulphate  (H4N)2  SO4— 

Is  amorphous  with  K2SO4  and  highly  soluble  in  water — its 
preparation  from  gas  water  given  above. 

Nitrate  (H4N)  N03— 

Is  amorphous  with  KNO3  and  deliquesces  in  air — decom- 
posed by  heat  to  N20  and  H20. 

H4N  NO3  = N20  “I-  2H20. 

Highly  soluble  in  water,  forming  a cooling  mixture. 


Nitrite  (H4N)N02— 

White  crystalline  solid  prepared  by  action  of  AgN02  on 
(H4N)C1— 

AgN02  + (H4N)C1  — AgCl  + (H4N)N02. 

Its  decomposition  by  heat  yields  pure  N — 

(H4N)N02  = N2  + 2 H20. 

Ammonium=Sodium=Phosphate  HNaNH4P04 ' 4H20 — 

“Microcosmic  salt”  is  formed  in  stale  urine — may  be  pre- 
pared from  HNa2P04  + (H4N)C1— 

HNa2P04  + H4NC1  = HNaNH4P04  + NaCl. 

Carbonate — 

White  or  transparent,  hard  mass  giving  off  H3N  to  the  air — 
formed  by  action  of  CaCC>3  on  (H4N)2S04 — 

(H4N)2S04  + CaC03  = CaS04  + (H4N)2C03. 

Commercial  salt  contains  also  the  bicarbonate,  HNH4CO3, 
and  the  carbamate,  NH2(NH4)C02,  which  are  converted  to 
(H4N)2C03  by  solution  in  (H4N)0H. 

Thiocyanate  (H4N)CNS— 

Prepared  by  action  of  CS2  on  H3N. 

CS2  + 4H3N  — H4NCNS  + (H4N)2S. 

Decomposed  by  heat  to  thio-urea. 

po-NHa 

Lb"NH2 

Sulphydrate  (H4N)SH — 

Prepared  by  running  H2S  into  ammonia  water. 

(H4N)0H  + H2S  = (H4N)HS  + h2o. 

H4N0H  will  now  change  it  to  (H4N)2S. 

(H4N)HS  + H4N0H  = (H4N)2S  + h2o. 

The  hydrosulphide  dissolves  S,  forming  the  yellow  ammonium 
sulphide,  which  contains  the  polysulphides  and  is  valuable  as  a 
reagent — (H4N)2S  decomposes  gradually. 

(H4N)2S  + H20  + 0 = 2(H4N)0H  + s. 

While  any  (H4N)2S  remains,  the  S is  dissolved,  but  finally  pre- 
cipitates. 


LITHIUM. 


History — 

Discovered  in  1017  by  Arfoedsen  in  “Petalite.” 


Occurrence — 


In  small  amounts  quite  widely  distributed — chiefly  as  a 
compound  silicate  in  lepidolite  or  “lithia  mica,”  found  also  in 
some  mineral  springs,  and  in  ashes  of  many  plants,  notably  to- 
bacco and  the  beet. 


Preparation — 

Metal  obtained  by  electroUses  of  the  chloride. 

Properties — 

Lightest  of  all  metals  (sp.  gr.  .59) — Oxidizes  without  melt- 
ing in  H2O — Most  Li  salts  are  soluble  in  water  and  give  a pur- 
ple red  color  to  the  flame.  Li  characterized  by  the  small  solu- 
biIit\T  of  Li3P04  also  of  LiCOs  (.75  in  100).  Otherwise  analo- 
gous to  Na  and  K. 

CESIUM— Cs— (sky-blue) . RUBIDIUM— Rb— (dark  red) . 

At.  Wt.  132.9.  Val.  I.  At.  Wt.  85.4.  Val.  I. 

Rare  metals  named  from  color  of  lines  they  give  in  spectro- 
scope, by  means  of  which  they  were  in  1860  discovered  by  Bun- 
son  and  Kirchhoff.  Both  are  perfect  analogues  of  K,  and 
though  rare  are  widely  distributed.  They  frequently  accom- 
pany K in  mineral  springs  and  plant  ashes,  and  occur  in  larger 
quantities  in  lepidolite  (Rb.  5%)  and  pollucite  (Cs.  30%). 


METALS  OF  THE  ALKALIES. 


Symbol: 

Li. 

h4n. 

Na. 

K. 

At.  Wt.: 

7 

18 

23 

39.1 

MsSbO*: 

soluble 

insoluble 

insoluble 

soluble 

MHC4.H4.Oe: 

soluble 

soluble 

soluble 

insoluble 

MoPtCh: 

soluble 

insoluble 

soluble 

insoluble 

MNOs: 

deliquescent 

deliquescent 

deliquescent 

permanent 

M2CO3: 

difficultly  sol. 

soluble 

efflorescent 

deliquescent 

I/3PO4: 

difficultly  sol. 

insoluble 

soluble 

soluble 

Salts  in  general: 

permanent 

volatile 

efflorescent 

deliquescent 

CALCIUM— Ca. 

At.  Wt.  40.  Val.  II. 


History — 

In  1800  Davy  obtained  the  metal  from  CaO. 


Occurrence — 

Widely  distributed  in  large  quantities — asCaO  inmost  min- 
erals— as  CaCOa  in  limestone,  chalk,  marble  -as  CaSC>4  in  gyp- 
sum, alabaster,  etc. — as  CaSC>4  and  CaCoa  in  most  natural 
waters — likewise  contained  as  phosphate  and  carbonate  in 
shells  and  bones  of  animals,  and  all  ashes  of  plants. 

Preparation — 

Metal  best  obtained  by  electrolyses  of  CaCU. 

Properties — 

Yellow  ductile  metal — decomposes  water — heated  in  the  air 
Ca  burns  to  CaO. 

CALCIUM  OXIDES. 

Calcium  forms  two  oxides,  CaO  and  Ca02. 

Ca02  is  formed  by  heating  CaC03  in  streatn  of  0. 

The  more  important  is  CaO,  commonly  called  “lime”  or 
“quicklime.”  Pure  CaO  is  a white,  infusible  solid,  best  prepared 
b}^  heating  marble  or  Iceland  spar  (CaCOs). 

heated 

CaCOs  = CaO  + C02. 

When  exposed  to  moist  air  CaO  again  takes  up  water,  forming 
Ca(OH)2,  which  then  unites  with  CO2,  forming  the  carbonate. 
By  reason  of  its  infusibilitv  CaO  gives  with  the  compound 
blow-pipe  the  intense  calcium  light. 

CALCIUM  HYDRATE— Ca(OH)2. 

Preparation — 

“Slacked  lime”  is  formed  b\^  * 

(1)  Adding  H2O  to  CaO. 

CaO  + H20  =='  Ca(OH)2. 

Reaction  is  accompanied  by  considerable  heat,  sufficient  to 
explode  powder  or  char  wood. 

(2)  Treating  CaCL  with  KOH. 

CaCl2  + 2KOH  = Ca(OH)2  + 2KC1. 

Properties — 

Soft  white  powder — attracts  CO2  of  the  air  to  form  CaCOa 
— at  red  heat  the  hydrate  breaks  up  to  CaO  + H2O.  Ca(OH)2 
is  slightly  soluble  in  cold,  and  less  in  hot  water,  whence  a sat- 


tira ted  cold  solution  becomes  cloudy  when  heated.  When  CaQ 
is  added  to  water  in  excess,  a portion  of  the  hydrate  goes  into 
solution,  the  remainder  is  precipitated.  The  clear  supernatant 
liquid  is  used  as  lime  water,  the  bulky  precipitace  as  “milk  of 
lime.” 

Lime  water  has  an  alkaline  reaction — unites  easity  with 
CO2 — this  is  used  as  a test  for  CO2  in  the  air — crystallized 
Ca(0H)2  may  be  obtained  from  lime  water — the  solubility  of 
Ca(0H)2  is  increased  by  presence  of  sugar  and  diminished  by 
alkalies. 

Slacked  lime  is  used  largely  in  ordinary  mortar,  a mixture 
of  lime,  water  and  sand,  in  which  Ca(0H)2  forms  CaC03  from 
the  CO2  of  air,  and  likewise  forms  with  the  sand  a calcium  sili- 
cate. Hydraulic  mortar  is  produced  by  igniting  limestone 
with  aluminium  silicate,  the  composition  being  chiefty  Ca  and 
A1  silicates. 

Ca(0H)2  is  cheapest  of  all  bases  and  hence  used  wherever 
possible  in  preparing  other  hydrates. 

CALCIUM  CHLORIDE  CaCl2. 

Made  by  treating  CaC03  with  HC1. 

CaC03  + 2HC1  = CaCl2  + C02  + H20. 

Some  Ca(0H)2  must  be  added  to  precipitate  Fe  or  Mn  present 
as  impurities. — Chloride  obtained  by  evaporation  has  formula 
CaCL  ‘ 6H2O — is  very  soluble  inH20and  is  used  in  freezing  mix- 
tures.— Anhydrous  CaCL  is  a white  crystalline  solid  used  to 
dry  gases. 

Bleaching  Powder  CaOCL — 

Treat  slaked  lime  with  Cl  gas. 

Ca(OH)2  + 2C12  = CaOCl2  + CaCl2  + 2H20. 

Formula  of  bleaching  powder  is  undecided,  but  is  approximate- 
ly CaClOCl  (CaOCL) — commercially  called  “chloride  of  lime,’ r 
“chlorinated  lime,”  “chemick.” 

Properties — 

White  porous  powder — odor  of  chlorine — CO2  of  the  air 
decomposes  it,  freeing  HC10 — by  action  of  strong  acids  Cl  is  set 
free. 


-60 


Ca(C10)2  + 2HC1  — CaCl2  + 2HC10. 

HC10  + HC1  = Cl2  + H20. 

A weak  acid  as  CO2  sets  free  HC10 — its  strong  bleaching 
and  disinfection  powers  are  due  to  these  reactions — in  bleaching, 
an  alkaline  bath,  also  an  anti-chlor  bath  should  follow. 

anti-chlor 

2Na2S203  + Cl4  = 4NaCl  + 3S02  + S. 


CALCIUM  SULPHATE  CaS042H20. 

Occurrence — 

Anhydrous  forms  occur  as  “gypsum”  of  which  alabaster  is 
one  variet}r — natural  waters  contain  CaS04  in  solution. 

Properties — 

Not  easily  soluble  in  hot  or  cold  water — presence  of  H4N  or 
Na  salts  or  acids,  markedly  increases  its  solubility.  Insoluble 
in  C2H6O  (alcohol).  Heated  to  100°  it  loses  most  of  its  crystal 
water,  forming  plaster  of  Paris,  an  amorphous  powder — 
this  will  again  unite  with  water  to  form  the  crystalline  variety 
— the  “setting”  is  due  to  crystallization,  which  can  be  slowed 
by  Na2B407,  and  hastened  by  K2S04.  If  heated  above  200°  it 
becomes  too  dense  to  set. 

CALCIUM  PHOSPHATES. 

Tertiary  or  Neutral  Ca3(P04)2 — 

Most  important  natural  occurrence  is  in  bones  (see  P). 

May  be  obtained  from  CaCl2  and  NasP04. 

3CaCl2  + 2Na3P04  = Ca3(P04)2  + 6NaCl. 

Or  from  the  di-sodic  phosphate  (Na2HP04)  in  alkaline  solution. 
2HNa2P04  + 3CaCI2  + 2NH3  = Ca3(P04)2  + 4NaCl  + 2(H4N)C1. 
When  fresh  is  a gelatinous  mass  which  dries  to  a white 
powder  highly  insoluble  in  water  but  easily  converted  to  the 
soluble  forms. 

Secondary  CaHP04— 

A white  crystalline  salt  formed  by  treating  Na2HP04  with 
CaClo. 

Na2HP04  + CaCl2  = CaHP04  + 2NaCl. 


61 


Primary  or  Acid  H4Ca(P04)2— 

White  crystalline  salt — soluble  in  H2O  with  acid  reaction — 
prepared  by  dissolving  HCaPC>4  in  HNO3  and  evaporating. 
2HCaP04  + 2HN03  = CaH4(P04)2  + Ca(N03)2. 

Largely  used  as  a fertilizer. 

Hypophosphite — 

White  crystalline  salt — obtained  by  action  of  P on  Ca(OH)2 
— (see  preparation  of  H3P). 

Calcium  Nitrate  Ca(N03)2~ 

A white  deliquesent  salt  of  little  importance. 

Calcium  Carbonate  CaC03 — 

Widely  distributed  as  limestone,  chalk,  marble,  etc.,  purest 
in  marble  and  Iceland  spar — occurs  likewise  in  forms  of  plant 
and  animal  life — insoluble  in  pure  water,  it  dissolves  somewhat 
in  H2CO3,  hence  found  in  all  natural  waters — Heated  it  readily 
yields  CO2  and  CaO  and  is  thechief  source  of  Ca  salts  and  CO2. 

Calcium  Sulphide-^ 

Ca  forms  polysulphides  of  which  CaS  is  the  most  important 
— Calcium  sulphide  is  a yellowish  mass,  formed  by  heating 
CaSC>4  with  C — chiefly  notable  from  it  phosphoresence,  hence 
used  in 'preparation  of  luminous  paints. 


REACTIONS  OF  CALCIUM  SALTS. 


NaOH  and  KOH: 
HiNOH: 

HzCrCU: 

H2SO4.: 

(if2iV)2p204: 

Flame: 

Blowpipe: 


precipitate  Ca  salts  as  Ca(0H)2. 

gives  no  precipitate. 

gives  no  precipitate  (compare  Ba). 

gives  no  precipitate  (compare  Ba  and  Sr). 

precipitates  in  alkaline  solutions. 

P insoluble  in  H(C2Hs02) — soluble  in  mineral  acids, 
reaction  is  a red  color  when  salt  is  moistened  with  HC1. 
Ca  salts  give  a luminous  residue. 


STRONTIUM— Sr. 

At.  Wt.  87.5.  Val.  II. 

The  earth  Strontia  discovered  in  1793 — the  metal  by  Davy 
in  1808  b}^  electrolysis  of  the  hydrate. 

A rare  metal  found  as  the  carbonate  and  sulphate. 


Properties — 

In  general  analogous  to  Ca. 

Strontium  Oxide  SrO — 

Prepared  like  CaO  from  the  carbonate. 

Strontium  hydrate  Sr(OH)2 — 

A strong  base  resembling  Ca(OH)2  but  more  soluble  in 
water.  Forms  salts  like  Ca(OH)2. 

REACTIONS  OF  Sr  SALTS. 

NaOH  and  KOH:  gives  precipititate  of  Sr(OH)2  only  in  concentrated  solution. 
H2SO4:  or  soluble  sulphates,  precipitate  SrSCb 

ATa2C03:  precipitates  SrC03. 

Na2HPO±:  precipitates  SrHP04. 

The  only  important  salt  is  the  nitrate,  much  used  in  fire- 
works. 

BARIUM. 

At.  Wt.  137.  Val.  II. 

H istory — 

Metal  obtained  by  Davy  1808. 

Occurrence — 

Chiefly  as  the  carbonate  (witherite)  and  the  sulphate  (bar- 
ite). 

Preparation — 

Best  obtained  by  electrolysis  of  BaCl2.  C will  not  reduce 
the  metal  from  BaCOs,  as  it  does  with  the  alkali  carbonates. 

Properties — 

Analogous  to  Sr  and  Ca,  but  a stronger  base. 

Barium  Oxide  BaO — 

Prepared  (1)  by  igniting  Ba(NOs)2. 

Ba(N03)2  = BaO  + 2N02  + O'. 

(2)  Heating  BaCOs  with  C. 

BaCOs  + c = BaO  + 2C0. 

Properties  are  similar  to  those  of  CaO. 

Barium  Peroxide  Ba02 — 

Made  by  heating  BaO  in  stream  of  O. 

BaO  + O — Ba02. 

If  strongly  heated,  Ba02  yields  its  O. 

Ba02  = BaO  H-  0. 


63 


This  is  one  source  of  commercial  0.  Ba02  is  used  also  in 
preparing  Ozone  and  Hydrogen  peroxide. 

Barium  Hydrate  Ba(0H)2~r 

BaO  + H20  = Ba(OH)2  + heat. 

Ba(OH)2  crystallizes  with  8 mols.  of  water,  Ba(0H)2'8H20 
— with  water  forms  an  aquous  solution  called  “Baryta  water” 
which  is  a strong  base  and  like  Ca(0H)2  rapidly  absorbs  CO2 
from  air — is  considered  a somewhat  better  reagent  for  carbonic 
acid. 

REACTIONS  OF  Ba  SALTS. 


Barium  salts  similar  to  Ca  salts  but  more  poisonous. 


NaOH : 
H^NOH : 
H2SO4, : 
K2CrO 4 : 
Na2C03 : 


precipitates  Ba(OH)2  only  in  strong  solutions, 
gives  no  precipitate. 

precipitates  the  most  dilute  solutions,  j Differ  from  Ca  and  Sr 
precipitates  insoluble  BaC03. 


The  sulphides  of  Ba,  Sr,  and  Ca  are  all  phosphorescent. 
Ba(N0s)2  gives  a green  fire  in  pyrotechnics. 


METALS  OF  THE  ALKALINE  EARTHS. 


Symbol: 

Ca. 

Sr. 

Ba. 

At.  Wt: 

40 

87.5 

137 

M(OH)2: 

insoluble 

medium 

soluble 

M2C03: 

insoluble 

insoluble 

insoluble 

MCh: 

soluble 

soluble 

soluble 

soluble  in  C2HrO 

soluble  in  C2HeO 

insoluble  in  C2HeO 

deliquescent 

permanent 

permanent 

M(N03)2: 

deliquescent 

permanent 

permanent 

M2SO±: 

1 : 400 

1 : 6895 

1 : 685,000 

M2CrO±: 

soluble 

soluble 

insoluble 

MAGNESIUM— Mg. 

At.  Wt.  24.  Val.  II. 

History — 

In  1750  Black  distinguished  magnesia,  MgO,  from  lime, 
CaO.  I11  1808  Davy  showed  that  it  was  the  oxide  of  a metal. 
Later  Wohler  obtained  the  pure  metal. 

Occurrence — 

Compounds  less  abundant  than  those  of  Ca  but  as  widely 
distributed — contained  in  most  soils  and  all  plant  ashes— in 


/ 


64 

large  quantity  in  Hornblende  and  meerschaum  (silicates)  and 
dolomite  (carbonate  of  Mg  and  Ca). 

Preparation  — 

Metallic  Mg  prepared  by  electrolysis  of  the  chloride  heated 
to  fusion,  or  heating  the  chloride  with  Na. 

Properties — 

Lustrous  silver  white  metal — medium  hard,  ductile,  and  mal- 
leable. At  ordinary  temperature  Mg  gradually  oxidizes  on  the 
surface— ignited  burns  with  an  intense  light.  Mg  is  easily  at- 
tacked by  dilute  acids— slightly  by  the  haloids — does  not  de- 
compose H2O,  as  Mg(0H)2  is  insoluble. 

Uses — 

Metallic  Mg  is  made  chiefly  to  burn , and  comes  next  to  lime  in 
strength  of  light — used  largely  for  “flashlights”  in  photogra- 
phy— forms  many  alloys,  the  one  with  Zn  is  often  used  in  place  of 
pure  Mg  as  it  burns  with  an  equally  bright  light. 

Magnesium  Oxide  MgO— “Magnesia”— 

A white,  very  light  powder,  prepared  by  heating  magnesia 
alba  (MgC03*  Mg(0H)2).  Commercial  magnesia  contains 
traces  of  Si02,  FeO  and  CO2. — MgO  forms  with  H2O  an  insol- 
uble hydrate  Mg(0H)2. 

Magnesium  Hydrate  Mg(0H)2-^ 

Formed  as  above,  or  by  treating  any  Mg  salt  with  NaOH. 
Mg(0H)2  is  a white  amorphous  powder,  attracting  CO2 
from  the  air — a strong  base,  almost  insoluble  in  water,  but  sol- 
' uble  in  (H4N)C1 — is  used  withFe2(S04)3as  antidote  for  arsenic. 

riagnesium  Chloride  MgCl2 — 

A deliquescent  salt,  occurring  in  sea  water  and  mineral 
springs — prepared  from  MgC03  and  HC1. 

MgCOs  + 2HC1  =*  MgCl2  + C02  + H20. 

Heated,  the  deliquescent  crystals  yield  water  and  break  as 
below. 

MgCl2  + H20  = MgO  + 2HC1. 

If  H4NCI  should  be  present  we  would  obtain  anhydrous 
MgCl2 — 

The  commercial  salt  occurs  as  a bi-product  in  making  KC1 
and  is  much  used  in  the  arts. 


65 


Magnesium  Sulphate  MgSCLThLO — 

“Epsom  salts”  occur  in  sea-water  and  mineral  springs, 
especially  at  Stassfurt.  Pure  MgS04  forms  in  rhombic  efflores- 
cent crystals,  but  usually  contains  some  MgCl2.  MgSC>4  is  high- 
ly soluble  in  water  with  bitter  saline  taste — insoluble  in  alcohol. 

Used- 

In  medicine — also  to  weight  cotton  cloth— a saturated  solu- 
tion with  dextrine,  will  give  a crystalized  surface  on  glass. 

Magnesium  Carbonate  MgCC>3 — 

Occurs  in  compact  masses  as  magnesite,  also  with  CaC03 
as  dolomite. 

The  “magnesia  alba”  of  medicine  is  a basic  carbonate 
formed  by  treating  an  Mg  salt  with  NaOH.  Some  CO2  escapes 
and  a white  precipitate  falls,  which  when  dried  at  a low  tem- 
perature has  the  formula  Mg(OH)2' 4MgC03  + 4H2O. 

REACTIONS  OF  Mg  SALTS. 

precipitates  Mg(OH)2. 

“ “ incompletely;  with  (H4N)C1  no  precip. 

Mg(0H)2-MgC03. 

“ “ “ incompletely;  never  with  (H4N)  Cl 

no  precipitate ; compare  Ca. 

precipitate  MgHP(>4,  which  with  H4NCI  changes  to 

MgH4NP04 

BERYLLIUM. 

At.  Wt.  9.  Val  II. 

A rare  metal  found  native  only  as  the  oxide — common  form 
is  Beryl  (3  Be02Al203  * 6Si02) — when  green  it  is  the  emerald — 
Salts  resemble  those  of  Mg. 

ZINC— Zn. 

At.  Wt.  64.9.  Yal.  II. 

History — 

Brass  or  bronze  was  known  to  the  ancients  as  evidenced  in 
old  coins — Zn  was  first  recognized  as  a peculiar  metal  by  Para- 
celsus who  introduced  the  name,  Zinc. 

Occurrence — 

Is  comparatively  rare — found  native  only  in  Australia — 
chief  ores  are  “Smithsonite”  (ZnCC>3)  and  zinc  blende  (ZnS)— 
Qres  usually  accompanied  by  Cadmium. 


Na(OH): 
(H±N)OH : 
Na2COa\ 

(IL  N)2COs: 

(H4N)2C20±: 

Na2HPO±: 


G6- 


Extraction  from  ores — 

By  roasting  in  the  air  and  then  igniting  the  resulting  oxide 
with  C — 

ZnO  + C = Zn  + CO. 

Zinc  is  quite  volatile  and  hence  the  crucibles  must  be  connected 
with  a condenser — As  in  the  case  of  S the  Zn  vapor  condenses 
at  first  to  a fine  dust  called  Zinc  dust,  which  contains 
Zinc,  generally  cadmium,  and  all  the  volatile  impurities. — The 
commercial  Zinc  which  forms  after  the  dust,  contains  also 
many  impurities  especially  Pb,  Fe,  C,  sometimes  S and  Cd,  and 
small  amounts  of  As  and  Sb — it  is  purified  by  repeated  distilla- 
tion. 

Detection — 

Most  characteristic  compound  is  ZnS— before  the  blow-pipe 
it  gives  with  Cu(N0s)2  a green  color  known  as  Rinman’s  green 
— the  white  oxide  becoming  yellow  on  ignition,  is  also  charac- 
teristic. 

Properties — 

Blue  white,  crystalline  metal — physical  properties  vary  with 
temperature — brittle  at  ordinary  temperature — at  100-150°  can 
be  welded  or  drawn — 205°  breaks  under  the  hammer — melted 
and  poured  into  cold  water  forms  granulated  Zn — not  affected 
by  air  till  strongly  heated  when  it  burns  with  strong  blue  flame 
to  ZnO  — pure  Zn  is  slightly,  commercial  Zn  is  easily,  attacked  by 
acids  and  alkalis. 

Used- 

In  galvanizing  sheet-iron — forming  alloys,  especially  brass 
— in  galvanic  batteries — with  acids  and  alkaline  hydrates  as 
reducing  agent. 

Zinc  Oxide  ZnO— 

White  powder,  which  becomes  yellow  with  heat.  Prepared 
as  “Zinc  white,’ ’ a pigment  made  by  distilling  metallic  Zn— pre- 
pared for  medical  use  by  igniting  the  carbonate — for  dental  use 
by  igniting  the  nitrate,  the  last  two  methods  avoid  As — ZnO  oc- 
curs in  nature  as  Zincite. 


67 


Zinc  Hydrate  Zn(0H)2 — 

White  powder — formed  by  treating  Zn  salt  with  KOH,  and 
soluble  in  excess  of  alkali. 

ZnCl2  + 2KOH  = Zn(OH)2  + 2KC1. 

Zn(OH)2  + 2K0H  = Zn02K2  = 2H20. 

The  same  K zincate  formed  when  metallic  Zn  is  treated  with 
KOH. 

Zn  + 2K0H  v=  Zn02K2  + 2H20. 

This  reaction  often  used  in  alkaline  reduction. 

Zinc  Chloride  ZnCl2 — 

1)  Heating  Zn  in  current  of  Cl  gas. 

2)  Dissolving  Zn  in  HC1. 

Zn  + 2HC1  =?  ZnCl2  + H2. 

White  deliquescent  mass — fuses  with  heat— vaporizes  at  red 
heat — at  high  heat  forms  ZnO  Hr  HC1  (compare  MgCU) — the 
ZnO  unites  with  unchanged  ZnCl2  to  form  basic  chloride — a 
concentrated  solution  of  ZnCU  + ZnO  hardens,  with  evolution 
of  heat  to  ZnOHCl,  which  is  used  as  a dental  filling.  ZnCU  is 
used  as  disinfectant,  antiseptic,  and  wood  preservative. 

ZnBr2  and  Znl2  are  analogous  to  ZnCl2. 

Zinc  Sulphate  ZnS04*7H20— 

“White  vitriol”  is  found  in  Zn  mines — prepared  (a)  by  oxi- 
dation of  blende  (ZnS). 

ZnS  + 04  = ZnS04. 

(b)  by  solution  of  Zn  in  H2SO4. 

Zn  + H2S04  = ZnS04  + H2. 

By  ignition  it  forms  the  oxide  ZnO. 

Used  in  the  arts  and  medicine,  in  battery  fluids,  etc. — used 
also  to  form  colors  by  igniting  with  metallic  salts— emetic. 

Zinc  Carbonate  ZnCOs — 

Formed  in  company  with  the  hydrate,  as  basic  Zn  carbonate, 
when  a Zn  salt  is  treated  with  a soluble  carbonate. 

ZnCl2  + Na2C03  = ZnC03  + 2NaCl. 

ZnCl2  + Na2C03  + H20  = Zn(OH)2  + 2NaCl  + C02. 

Notice  tendency  to  form  hydrates  also  in  Mg. 


68 


Phosphates — 

Most  important  are  the  basic  phosphates — when  metaphos- 
phoric  acid  is  mixed  with  ZnO  it  forms  a cement,  which  is  a 
basic  metaphosphate. 

ZnO  + HP03  = Zn(0H)P03. 

Sulphide  ZnS— 

Zinc  blende  in  nature  is  yellow — artificial  is  white — is  readily 
formed  from  Zn  dust  and  powdered  sulphur.  ZnS  is  soluble  in 
mineral  acids  and  alkalies — a little  is  formed  by  action  of  H2S 
on  ZnS(>4,  but  the  acid  freed  soon  dissolves  it. 

Zinc  Salts— 

Are  colorless  if  the  acid  is  colorless — normal  and  soluble  salts 
redden  litmus — Zn  salts  are  poisonous,  have  metallic  and  a 
stringent  taste  and  used  as  emetics. 

CADMIUM-Cd. 

At.  Wt.  112.  Val.  II. 

History — 

Discovered  in  1817  by  Hermann  and  Stromeyer  simultane- 
ously. 

Occurrence — 

One  of  rarer  metals — generally  accompanies  Zn  ores — Z11 
blendes  often  contain  from  .2  to  3 percent  Cd— the  rare  mineral 
“Greenockite”  is  CdS. 

Preparation — 

Cd  is  obtained  in  the  first  part  of  Zn  distillate,  i.  e.,  zinc 
dust,  and  is  purified  by  fractional  distillation,  as  it  is  more  vol- 
atile than  Zn. 

Properties— 

A white,  brilliant  metal — malleable  and  tenacious,  but  the 
smallest  trace  of  Zn  makes  it  brittle. — Cd  is  not  oxidized  till  at 
high  temperature,  when  it  burns  to  CdO. — Cd  is  soluble  in 
H2SO4,  HC1  and  best  in  HNO3. — A pure  Cd  salt  is  precipitated 
completely  by  H2S. — An  amalgam  of  Zn  + Cd  is  used  as  a dental 
filling. 

Oxide  CdO — 

The  only  oxide  of  Cd  is  formed  by  burning  the  metal  or  ig- 
niting the  nitrate  or  carbonate — the  hydrates  are  formed  by 
KOH  acting  on  a Cd  salt. 


G9 


Sulphate  CdSCU — 

The  sulphate  Cd(S04)2  is  the  most  common  salt  of  Cd  and 
is  anhydrous. 

Nitrate  Cd(N03)2-4H20— 

The  nitrate  Cd(N0s)2'4H20  is  ver}'  soluble  and  gives  the 
oxide  on  ignition. 

Chloride  CdCl2— 

The  chloride  CdCU  is  similar  to  ZnCl2  but  efflorescent. 

Cadmium  Sulphide  CdS — 

The  most  important  Cd  salt  occurs  native  a Greenockite — 
commonly  made  from  H2S  and  a Cd  salt — not  easily  from  the 
elements. — CdS  is  scarcely  attacked  by  dilute  acids,  but  decom- 
posed by  strong — forms  a good  yellow  pigment — differs  from 
AS2S3  in  volatility  and  solubility  in  alkalies. 

CADMIUM  SALTS. 

In  general  resemble  Zn  salts. 

NaOH  and  KOH:  precipitate  Cd(0H)2  insol.  in  excess  (note  difference  from  Zn). 


{H±N)OH: precipitates  Cd(0H)2  soluble  in  excess. 

Na2COs: “ insoluble  basic  carbonate. 

Na2HP04!: “ “ CdHP04. 

iVa2C204: “ “ CdC204. 

K2CrO±: does  not  precipitate  dilute  solutions  (difference  from  Zn). 

H2S: precipitates  yellow  CdS  insol.  in  acids  (difference  from  Zn). 

BaCOz ’■ precipitates  CdC03  completely  (difference  from  Zn). 

COPPER— Cu. 


At.  Wt.  67.5.  Val.  II  and  I (-ic  and  -ous  salts). 

History— 

Known  to  the  ancients  before  iron  and  used  for  arms — name 
from  Cyprus,  where  the  Greeks  and  Romans  got  it. 

Occurrence — 

Native  in  large  quantities — ores  are  generally  the  oxide  and 
sulphide,  especially  chalcopyrite  (CuFeS2). 

Extraction — 

Two  processes,  the  “English’ * and  Mansfield,  the  latter 
used  when  ores  are  abundant  but  poor.— In  English  process  the 
divided  CuFeS2  is  first  roasted  in  the  air — this  converts  it  par- 
tially to  the  oxide— then  the  ore  is  ignited  with  silica  fluxes  and 
carbon — the  iron  reduces  to  the  oxide  and  is  dissolved  in  the  slag 


70 


— bv  repeating  the  process  a blistered  mass  is  obtained  which 
contains  much  S — this  is  repeatedly  roasted,  melted  and  “poled” 
and  Cu  run  into  iron  moulds.  The  Mansfield  process  is  the 
same,  save  that  a blast  furnace  is  used  and  more  silicious  flux. 

Properties — 

Red  crystalline  metal — malleable — tenacious — weldable — 
conducts  heat  and  electricity — not  attacked  by  dilute  acids  save 
in  presence  of  air — dissolves  slightly  in  HC1,  more  in  hot  H2SO4. 
Cu  + 2H2SO4  = CuS04  + S02  2 H20. 

best  in  HNO3. 

3Cu  + 8HNO3  - 3Cu(N03)2  + 4H20  + 2N0. 

In  dry  air  Cu  is  unaffected — in  moist  air,  it  forms  CuO — with 
CO2  is  coated  with  CUCO3 — H2S  blackens  it,  forming  CuS.— Cu 
is  blue  green  when  melted,  and  expands  on  cooling — melted  Cu 
always  contains  gas  which  escapes  when  it  solidifies. — This  is 
called  “spitting”  and  prevents  casting. 

Pure  Cu  is  used  as  wire,  roofing,  etc. — All  casting  must  be 
made  from  alloys,  of  which  brass,  german  silver  and  the  various 
bronzes  are  examples. 

Copper  forms  four  oxides,  CU4O,  CU2O,  CuO,  Cu02,  of  which 
CU4O  is  of  least  importance. 

Copper  Suboxide  CU4O — 

An  olive  green  powder  formed  by  reducing  cupric  hydrate. 

Cuprous  Oxide  CU2O — 

In  nature  as  cuprite,  and  formed  when  cupric  hydrate  is  re- 
duced by  grape  sugar. 

2 Cu(0H)2  — 0 = Cu20  + 2 H20. 

Cuprous  hydrate  Cu2(OH)2  is  a yellow  precipitate  which 
loses  water  on  boiling  to  form  the  red  CU2O.  CU2O  is  used  to 
give  red  color  to  glass. 

CUPROUS  SALTS 

are  not  important  and  are  easily  converted  to  cupric  form — 
they  are  generally  colorless,  red  or  yellow. 

KOH: precipitates  the  yellow  hydrate  Cu2(OH)2. 

Na2COs:-..  precipitates  the  yellow  carbonate  CU2CO3. 

(HityOH:  dissolves  salts  colorless,  but  turns  blue  by  oxidation. 


Cupric  Oxide  CuO — 

Found  as  the  mineral  cuprite — may  be  formed  by  oxida- 
tion of  Cuor  by  heating  the  hydrate  or  nitrate — remarkable  for 
its  ease  of  reduction. — The  hydrate  Cu(OH)2  is  a strong  base 
which  is  changed  to  the  oxide  even  when  in  water. 

Copper  Peroxide  CuC>2 — 

The  peroxide  is  a yellow  brown  powder  formed  by  treating 
Cu(OH)2  with  H202.~ 

CUPRIC  SALTS. 

Generally  blue  or  green  when  crystallized — colorless  when 
anhydrous — color  the  flame  blue  or  green — tartaric  acid,  sugar 
and  many  organic  compounds  prevent  precipitation  of  Cu(0H)2. 
NaOH: precipitates  blue  Cu(0H)2  (save  as  above). 

H&NOH: ....  precipitates  blue  Cu(0H)2,  this  is  soluble  in  excess  of  reagent  and 
the  solution  thus  formed  dissolves  cellulose. 

NazCO^: ....  precipitates  basic  CUCO3. 

H2S: precipitates  black  CuS. 

K^Fe(CN)e  precipitates  red  Cu2Fe(CN)6- 

Zn,  Fe}  Pb:  precipitates  metallic  Cu  from  solutions. 

Cuprous  Chloride  CU2CI2 — 

The  only  cuprous  salt  of  importance— formed  from  CU2O 
and  HC1. 

Cu20  + 2HC1  = Cu2Cl2  + H20. 

A white  powder,  insoluble  in  water, — absorbs  carbon  monox- 
ide CO,  hence  is  used  in  analysis. 

Cupric  Chloride  CuCL- 

Formed  by  dissolving  the  oxide  or  carbonate  in  HC1. 

CuC03  + 2HC1  = CuCl3  + H2Q  + C02. 

Cupric  Nitrate  Cu(N0s)2— 

Formed  by  action  of  HNO3  on  CUCO3. 

CuC03  + 2HN03  = Cu(N03)2  + H20  + C02. 

Cupric  Sulphate  CuS04‘5H20— 

Found  native  in  mines,  where  CuS  is  oxidized — may  be  made 
by  dissolving  Cu  in  H2SO4 — CUSO4  is  soluble  in  three  parts 
cold  cr  one-half  part  hot  water — 

“Blue  Vitriol”  is  much  used  in  the  arts  in  pigments,  copper 
platirg,  galvanic  batteries,  etc. 


Copper  Arsenite  CuHAs03— 

A greenish  yellow  precipitate,  formed  when  cupric  sulphate 
is  added  to  potassium  arsenite— commonly  known  as  Scheele’s 
Grefn. 

“Paris  green”  is  also  a copper  arsenite,  mixed  with  copper 
acetate.  These  colors  are  very  poisonous  and  give  off  AsH3  in 
the  presence  of  organic  matter— colors  may  .be  removed  from 
fabrics  by  (H4N)0H. 

Cupric  Carbonate  CuC03 — 

The  neutral  salt  CuC03  is  not  known — when  Na2C03  is 
added  to  Cu  salt  the  basic  carbonate  is  precipitated  as  either  a 
blue  precipitate  Cu(0H)2  + 2CuC03  called  “azurite”  or  a green 
Cu(OH)2  * CuC03  called  “malachite.” 

Cupric  Ferrocyanide  Ct^FeCeNe — 

When  acetic  acid  and  potassium  ferrocyanide  are  added  to 
a Cu  salt,  a red  brown  precipitate  of  Cu2FeC6N6  is  formed, 
which  constitutes  the  test  for  Cu. 

Cupric  Sulphide  CuS— 

A black  compound  found  in  nature,  or  precipitated  by  H2S 
or  (H4N)2S  from  solution  of  Cu  salt.  Insoluble  in  dilute  acids 
—slowly  changes  in  moist  air  to  CuS04. 

ALLOYS  OF  COPPER. 

The  addition  of  Sn  or  Zn  to  Cu  prevents  “spitting”  and 
forms  an  alloy  harder  than  Cu,  more  durable,  and  of  better 
color — likewise  more  sonorous. 

Ancient  bronze  = Sn  + Cu. — Modern  = Cu  + Sn  + Zn. — 
German  silver  = Cu  + Zn  + Ni. — Brass  = Cu  + Zn,  in  different 
proportions  and  sometimes  contains  not  more  than  .5  per  cent 
of  Pb,  which  facilitates  working — a varity  of  brass  alloys  is 
produced  by  varying  the  amounts  of  Cu  and  Zn — the  later  in- 
creases hardness  and  fusibility,  but  decreases  specific  gravity 
and  malleability.  Aluminum  bronze  = A1  -f-  Cu  in  varying 
proportions. 

In  electro-plating,  the  article  is  immersed  in  a CuS04  bath 
—the  negative  pole  of  an  electric  battery  is  connected  with  the 
object— CuS04  decomposes  and  precipitates  copper  on  the 


object — in  electro-typing,  a mould  of  plaster  of  Paris  is  covered 
with  graphite,  and  electro-plated  as  above. 


SILYER-Ag. 

At.  Wt.  108.  Val.  I. 

Occurrence — 

Widely  distributed — often  native — the  chief  ore  is  the  sulphide 
Ag2S  which  occurs  sometimes  native,  but  generally  with  some 
other  sulphides,  especially  galenite  PbS.  From  argentiferous 
(silver  bearing)  lead  ores,  most  commercial  silver  is  obtained. 

Extraction — 

Three  processes  are  in  use:  (1)  (Pattinson  Process.)  Gal- 

ena, a mixture  of  Pb  and  Ag  sulphides,  is  strongly  heated  to 
melting — much  of  the  lead  separates  out  and  some  is  united  as 
PbAg  alloy — the  pure  lead  crystallizes  first  and  is  removed, 
leaving  a readily  fusible  alloy  of  PbAg — by  repetitions  an  alloy 
rich  in  Ag  is  obtained  which  is  then  cupelled.  2)  (Parkes  or  Zinc 
process).  The  molten  alloy  is  treated  withZn — the  alloy  of  zinc 
and  silver  thus  obtained  is  treated  with  superheated  steam, 
which  oxidizes  the  Zn,  leaving  Ag  unchanged.  3)  (Amalgama- 
tion process  ) Galena  is  roasted  with  NaCl,  forming  AgCl. 
H2O  and  Fe  reduce  it  to  Ag,  which  is  separated  by  amalga- 
mating with  Hg,  and  then  distilling. 

Properties — 

The  whitest  metal — harder  than  gold  and  softer  than  Cu — 
very  malleable  and  tenacious — at  high  temperature  forms  a 
light  blue  vapor — melted  Ag  absorbs  O from  the  air  and  con- 
tracts on  cooling — its  allo3Ts  with  Cu  do  not  contract  and 
polish  better  than  Ag — silver  is  not  oxydized  at  ordinary  tem- 
peratures but  O3  changes  it  to  the  oxide  Ag20 — Ag  is  not 
affected  by  alkalies  or  KNO3 — unites  with  Cl,  Br  and  I at  ordi- 
nary temperature  and  decomposes  H2S  forming  Ag2S,  which  is 
soluble  in  KCN, — Ag  is  soluble  in  HNO3,  and  H2SO4  cone.,  but 
AgNC>3  is  insoluble  in  strong  HNO3,  hence  dilute  HNO3  should 
be  used  in  testing  for  chlorides. 

Silver  has  three  oxides,  Ag40,  Ag20  and  AgO,  of  which  the 
first  is  of  least  importance. 


— —74 ' 


Silver  Suboxide  Ag40 — 

A brown  powder  formed  like  (JU2O  and  easily  converted  to 
Ag'jO  and  Ag2. 

Silver  Oxide  Ag20— 

The  most  important  oxide  of  silver  is  a dark  brown  precip- 
itate formed  when  KOH  is  added  to  AgNC>3—  any  Ag  hydrate 
formed  is  easily  decomposed,  leaving  Ag20,  which  acts  as  if  it 
were  Ag(OH) — a strong  base,  slightly  soluble  in  water  and  re- 
duced by  light,  heat  orH  at  100°.  Its  soluble  salts  are  poison- 
ous and  have  metallic  taste. 

Peroxide  AgO- 

Black  crystals  with  metallic  lustre  formed  by  passing  ozone 
over  Ag  or  Ag20. 

Silver  Chloride  AgCl — 

Found  in  nature  as  “Horn  silver,”  and  formed  when  HC1  is 
added  to  a soluble  Ag  salt — AgCl  forms  as  a white,  curdy  pre- 
cipitate insoluble  in  dilute  acids,  readily  soluble  in  (H4N)0H,. 
whence  it  c^stallizes  in  large  octohedrals — Fused  AgCl  sodifies- 
to  a horn-like  mass,  hence  the  name. 

Silver  Bromide  AgBr — 

Separates  from  silver  salts  on  addition  of  HBr  or  a soluble 
bromide — a bright  yellow  precipitate  similar  to  the  chloride  but 
less  soluble  in  (H4N)0H. 

Silver  Iodide  Agl — 

Yellow,  curdy  precipitate  formed  from  the  action  of  HI  ora 
soluble  iodide,  on  silver  salts  in  solution— insoluble  in  H4NOH. 

Photography — 

Certain  chemical  compounds  are  sensitive  to  sunlight  or 
other  chemically  active  rays,  (as  Mg  or  P light) — Prominent 
among  these  are  the  Argentic  haloids  (AgX),  which  form  vio- 
let changing  to  black  argentous  compounds  (Ag2X) — This  is 
the  basis  of  photography.  AgCl  is  the  most  sensitive  to  light, 
but  does  not  develop  well;  hence  the  “wet”  plate  is  sensitized 
by  a mixture  of  AgBr  and  Agl — The  “dry”  plate  with  AgBr. 
The  exposed  plate  is  treated  with  some  “developer”  like  pyro- 
gallic  acid,  which  reduces  the  argentous  haloid  (Ag2Br)  formed 
to  metallic  silver  and  dissolves  free  bromine — Metallic 


silver  unites  with  Argentic  Bromide,  forming  argentous,  which 
is  again  reduced  by  developer — Process  continues  till  sufficient 
‘ ‘molecular”  silver  is  formed  to  give  a distinct  image  on  the 
plate — After  sufficient  development  the  plate  is  transferred  to  a 
“fixer”  of  sodium  hyposulphite  (“hypo”),  which  reacts  with 
unchanged  silver  haloid,  freeing  Br  and  forming  the  soluble 
Na2S203  . Ag2S203. 

Silver  Sulphate  Ag2S04 — 

Obtained  by  dissolving  Ag  in  hot  H2SO4 — is  used  in  refining 
of  Ag — slightly  soluble  in  H2O. 

SILVER  NITRATE— AgN03. 

Preparation — 

“Lunar  caustic”  is  prepared  by  dissolving  Ag  in  HNO3. 

Properties — 

A white  crystalline  solid,  soluble  in  water,  ether  and 
alcohol — a brittle  substance,  but  less  so  if  accompanied  by 
AgCl,  which  is  generally  present  in  the  commercial  article — 
when  pure  is  unaffected  by  light  but  in  presence  of  organic 
matter  is  reduced  to  metallic  Ag. 

Silver  Fulminate  Ag2(CN)202 — 

Made  by  dissolving  AgNOs  in  C2H5OH  and  HNO3 — more 
explosive  than  the  Hg  salt. 

Cyanide  AgCN— 

AgCN  is  formed  when  HCN  is  added  to  solution  of  AgNOs 
— used  in  electro-plating. 

Carbonate  Ag2C03— 

Grayish  white  precipitate,  formed  when  Na2C03  is  added 
to  a silver  salt. 

Chromate  Ag2Cr04 — 

Ag2Cr04  is  precipitated  by  K2Cr04  from  soluble  Ag  salt — a 
yellow  precipitate,  which  is  an  indicator  in  analysis,  where  Ag 
salt  is  used. 

Silver  Sulphide  Ag2S — 

Occurs  native  as.  Argentite,  or  precipitated  by  H2S  or 
(H4N)2S  from  Ag  salts — a black  precipitate  used  in  silvering 
mirrors. 


76 


Alloys — 

Silver  is  generally  used  as  an  alloy  with  copper — In  coin 
90%  is  Ag,  in  silverware  about  75%,  “sterling”  silver  92%. 


SILVER  SALTS. 

NaOH gives  precipitate  of  Ag20. 

(H^N)  OH gives  precipitate  of  AgoO,  soluble  in  excess;  (in  presence  of 

free  acid  no  precipitate). 

Na2C03 gives  precipitate  of  3'ellowish  white  Ag2C03. 

NazHPO± gives  precipitate  of  yellowish  Ag2HP04- 

H2S  and  (//4AO2S  give  precipitate  of  black  Ag2S. 

HC1 gives  precipitate  of  white  curdy  AgCl  [sol.  in  (H4N)0H, 

KCN,  and  Na2S203 — insoluble  in  acids.] 

KI gives  precipitate  Agl  [insoluble  in  (H4N)0H,  but  soluble  in 

Na2S203-] 

FeSO± gives  precipitate  of  metallic  Ag. 

Tannic  Acid.........  gives  precipitate  of  metallic  Ag. 

Zn  and  Fe gives  precipitate  of  metallic  Ag. 

All  Ag  salts  are  anhydrous. 


MERCURY-Hg. 

At.  Wt.  200.  Val.  I and  II. 

History — 

Mercury  has  been  known  since  the  earliest  times. 

Occurrence — 

Not  widely  distributed  nor  in  large  quantities— comes  most- 
ly from  Spain  and  California — native  in  small  quantity — chief 
ore  in  Cinnabar  HgS— Rare  ores  are  “Amalgam”  AgHg  and 
and  “Horn  mercury”  2Hg2CL. 

Extraction — 

Native  Hg  is  filtered  through  leather — cinnabar  is  heated 
with  O. 

HgS  + 2 0 = Hg  + S02. 

Commercial  is  sometimes  pure,  but  often  contains  Pb,  Sn,  Bi 
and  Cu  dissolved — Pure  Hg  can  be  made  from 
HgCls  + Fe  = Hg  + FeCl2. 

Properties — 

The  only  liquid  metal— bluish  white  color,  metallic  lustre — 
at  40°  crystallizes  in  octohedra — evaporates  at  ordinary  tem- 
peratures— Hg  resembles  AuandPtin  its  small  affinity  for  0 and 
large  affinity  for  Cl— requires  high  temperature  to  form  the 


•77 


oxide,  HgO,  and  at  still  greater  heat  yields  it  up  again  (com- 
pare  Ba02) — Ozone  and  chlorine  attack  it  at  normal  tempera- 
ture, but  boiling  HC1  is  harmless — Hg  is  easil}’  divided  bv  fats 
and  chalk. 

Used  in  manufacture  of  thermometers  and  barometers — tin 
amalgam  for  mirrors — extracting  gold  and  silver  from  ores — 
useful  also  in  electric  connections. 

Mercurous  Oxide  Hg20 — 

Mercury  forms  two  oxides,  Hg20  and  HgO — Hg20,  cor- 
responding to  Ag20,  is  a heavy  black  powder  formed  from 
Hg2(N03)2  and  KOH— easily  decomposed  by  light  to  HgO  + O 
also  by  heat  first  to  HgO  + O,  then  to  Hg  +O2. — HC1  converts 
it  to  Itg2Cl2  ; Hg20  is  the  basis  of  mercurous  salts. 

MERCUROUS  SALTS. 


KOH : precipitates  black  Hg20. 

: precipitates  black  Hg20  containing  NH3. 

Na2CO%\ precipitates  yellow  unstable  Hg2C03  (which  changes  to. 

the  black  oxide  when  heated). 

H2S  and  (H4,N)2S:  precipitates  black  Hg2S. 

KI : precipitates  green  Hg2I2. 

HC1 : precipitates  white  Hg2 Cl 2. 

SnCl2 : precipitates  white  Hg2Cl2  (afterwards  black  Hg). 

Cu : precipitates  metallic  Hg. 


Mercurous  salts  are  mostly  insoluble — have  a metallic  taste 
and  act  less  violently  than  mercuric  salts  (Hg). 

riercuric  Oxide  HgO— 

Is  obtained  as  the  red  or  yellow  allotropic  modification  ac- 
cording to  preparation — Red  oxide  by  heating  Hg(N03)2. 
Hg(N03)2  — HgO  + 2NO. 

Yellow  oxide  in  the  wet  way  from  HgCL  and  NaOH. 

HgCl2  + 2 NaOH  = HgO  + 2NaCl  + H20. 

With  heat  HgO  changes  to  the  black  oxide  Hg20,  but  recovers 
its  oxygen  on  cooling — The  yellow  and  red  differ  slightly  in 
chemical  action  ; for  example,  the  yellow  is  at  once  converted 
to  the  oxalate  by  H2C2O4,  but  the  red  must  be  long  heated — both 
are  strong  oxidizing  agents  and  form  the  basis  of  mercuric  salts. 


REACTIONS  OF  MERCURIC  SALTS. 


KOH precipitates  the  yellow  oxide  HgO. 

H^NOH...  gives  a “white  precipitate”  in  very  dilute  solutions. 

KI gives  a yellow  precipitate  of  Hgl,  which  turns  red. 

H2S in  large  amounts  gives  black  HgS. 

H2S in  small  amounts  gives  yellow  precipitates  HgX2  . HgS.* 

Cu is  covered  with  a deposit  of  metallic  Hg. 


Mercurous  Chloride  Hg2Cl2— 

“Calomel”  is  prepared  in  (1)  Dry  way 

Hg  + HgCl2  = Hg2Cl2  (Crystalline). 

(2)  Wet  way 

White  powder. 

Hg2(N03)2  + 2HC1  = Hg2Cl2  + 2HN02. 

Hg2Cl2  sublimes  without  melting — somewhat  less  volatile  than 
HgCL — insoluble  in  H2O,  ether  and  alcohol — not  easily 
attacked  by  acids,  but  boiled  with  HC1  is  converted  to  HgCL. 

Mercuric  Chloride  HgCL — 

“Corrosive  sublimate”  is  formed  (1)  from  HgO  and  HC1 — 
HgO  + 2HC1  = HgCl2  + H20. 

(2)  HgS04  + 2NaCl  = HgCl2  + Na2S04. 

HgCL  is  a highlv  volatile  substance — soluble  in  cold  water, 
better  in  hot — soluble  also  in  alcohol  or  ether — aqueous  solu- 
tion reacts  acid  and  in  light  decomposes  to  Hg2CL — when 
boiled  some  of  HgCL  is  volatilized — mercuric  chloride  is  precip- 
itated by  albumenoid  substances,  hence  these  are  the  best  anti- 
dotes. 

Mercurous  Nitrate  Hg2(N0s)2 — 

Made  by  action  of  dilute  HNO3  on  excess  of  Hg — unless 
metallic  Hg  is  present  it  gradually  changes  to  Hg(N0s)2. 

Mercuric  Nitrate  Hg(N0s)2 — 

Solution  may  be  obtained  by  dissolving  Hg,  or  HgO  in  ex- 
cess of  strong  HNO3 ; acid  must  be  in  excess  or  basic  salt  will 
separate — the  basic  salt  is  reconverted  by  boiling  in  H2O. 

Mercuric  Cyanide  Hg(CN)2 — 

Chief  source  of  (CN)2  gas  and  is  prepared  from  HgO  and 
HCN. 

HgO  + 2HCN  Hg(CN)2  + H20. 


*X  stands  for  any  non-metal. 


79 


Mercuric  Fulminate  HgC2N2(>2 — 

Analogous  to  the  Ag  salt  but  less  explosive — used  in  percus- 
sion caps. 

* 

Tlercuric  Sulphide  HgS — 

Occurs  in  Cinnabar  and  Vermillion— Is  produced  as  black 
amorphous  mass,  by  action  of  H2S  on  a mercuric  salt — If  this 
black  sulphide  is  heated  it  sublimes  as  a red  crystalline  mass. 

SUBSTITUTED  H4N  COMPOUNDS. 

When  H4NOH  is  poured  over  Hg2Cl2  it  blackens  with  the 
following  reaction. 

amid  o-m  ercurous-chloride 

Hg2Cl2  + 2NH3  = NH2Hg2Cl  + H4NCI. 

The  new  compound  is  regarded  as  ammonium  chloride  in  which 
two  hydrogens  are  replaced  by  Hg2.  This  one  of  several  substitu- 
tion products  formed  by  the  entrance  of  Hg  into  ammonium 
compounds. 

METALS  OF  THE  Cu  GROUP. 


Cu 

Ag 

Hg 

Specific  Gravity 

8.95 

10.57 

13.5*9 

Melting  Point 

1054° 

954°......... 

39° 

At.  Wt 

63.4 

107.6 

199.8 

Ag  resembles  somewhat  the  alkali  metals — group  chlorides 
with  symbols  MCI2  are  soluble  in  H2O. 

Chlorides  M2CI2  are  insoluble. 

Sulphides  are  black,  and  insoluble  in  acids  or  alkalies  but 
CuS  is  slightly  soluble  in  (H4N)2S. 

LEAD-Pb. 

At.  Wt.  208.  Val.  II. 

History — 

One  of  earliest  known  metals. 

Occurrence — 

Occasionally  native — chief  ore  is  Galena  PbS,  a very  com- 
mon and  widely  distributed  mineral. 

Extraction — 

(1)  By  heating  the  sulphide  with  Fe. 

PbS  + Fe  = FeS  + Pb. 


80 


(2)  By  roasting  until  the  sulphide  is  partially  converted  to 
PbO  and  PbS04. 

PbS  + 30- J‘b0  + S02. 

And  PbS  + 04  - PbS04. 

By  ignition,  the  lead  now  separates. 

PbS  + 2PbO  = 3Pb  + S02. 

And  PbS  + PbS04  = 2Pb  + 2S02. 

Pure  lead  is  obtained  by  heating  PbC204  with  C. 

Properties — 

Lustrous,  blue  white  metal — at  1,600°  gives  off  poisonous 
vapors — all  its  salts  are  poisonous — at  red  heat  is  easily  oxi- 
dized— at  normal  temperature,  a thin  coat  of  the  oxide  forms, 
which  protects  the  body  of  metal. — Pb  is  dissolved  by  HNO3 
but  not  by  HC1,  nor  dilute  H2SCL  because  of  insoluble  salts 
formed — From  its  nitrate,  Pb(N0s)2,  lead  is  precipitated  in 
metallic  form,  by  Zn,  forming  a lead  tree.  WhenPb  is  in  con- 
tact with  air  and  water  the  somewhat  soluble  hydrate 
Pb20(0H)2  is  formed.  This  is  especially  dangerous  where  pure 
H2O  or  rain  water  stands  in  lead  cisterns  and  pipes.  With 
hard  water  the  action  is  very  slight. 

Salts  and  organic  matter  modify,  CO2  and  CaH2( 003)2 
and  sulphates  diminish,  chlorides  and  nitrates  and  especially 
nitrogenous  organic  matter  increase,  corrosion.  Lead  in  con- 
tact with  wood  is  rapidly  corroded. 

Lead  Suboxide  Pb203 — 

Lead  forms  four  oxides,  Pb20,  PbO,  Pb203  and  Pb02. 

Pb20  is  the  gray  coating  on  metallic  Pb — it  decomposes  to 
Pb  + PbO. 

Lead  Oxide  PbO — 

Found  occasionally  in  nature — artificial  is  of  various  colors 
accordingto  preparation,  i.e.,  red,  yellow  and  white,  due  proba- 
bly to  crystallization. 

By  ignition  of  Pb  we  get  the  yellow  which  becomes  red 
when  rubbed — when  strongly  heated  becomes  yellowish  red, 
forming  “litharge”. — PbO  is  a strong  base  resembling  BaO  and 
SrO — it  absorbs  CO2  from  the  air,  forming  PbC03. 

Lead  oxides  are  much  used  in  lead  plaster,  glazing  earthen 
ware,  making  glass,  etc. 


81 


Lead  Trioxide  and  Dioxide  Pb2C>3,  Pb02 — 

Pb203isof  no  importance.  Pb02  is  formed  on  positive  pole 
in  the  electrolysis  of  a lead  salt — also  by  action  of  Cl  on  PbC03 
— a dark  brown  powder,  which  conducts  itself  somewhat  like 
Mn02,  as  for  example,  when  heated  with  HC1,  chlorine  gas  is 
formed. 

iTinium  Pb3C>4 — 

“Red  lead”  is  a mixture  of  PbO  and  Pb203 — important  as 
a pigment,  and  made  b}T  oxidizing  Pb. 

Lead  Hydrate  Pb20(0H)2— 

When  a Pb  Salt  is  treated  with  KOH  it  does  not  form 
Pb(OH)2,  but  a basic  hydrate  of  approximately  Pb20(0H)2 — 
In  presence  of  strong  bases,  Pb20(0H)2  acts  as  an  acid,  hence 
is  soluble  in  KOH. 

Lead  Chloride  PbCl2— 

A white  precipitate  formed  when  HC1,  or  a soluble  chlo- 
ride, is  added  to  a cold  solution  of  Pb  salt.  Soluble  in  hot 
water,  but  is  largely  deposited  in  needle  shaped  crystals  when 
cool — melted  PbCl2  solidifies  to  a horn  like  mass. 

Bromide  PbBr2 — Iodide  Pbl2 — 

PbBr2  is  very  similar  to  the  chloride — chief  characteristic  of 
Pbl2  is  that  it  crystallizes  from  its  hot  water  solution,  in  yellow 
lustrous  laminte. 

Lead  Sulphate  Pb(S04) — 

Found  in  nature,  or  prepared  by  adding  a soluble  sulphate 
to  a lead  salt — almost  insoluble  in  water  and  dilute  H2SC>4 — 
soluble  in  HN03  and  (H4N)  C2H302. 

Nitrate  Pb(N03)2— 

The  most  common  soluble  salt  of  PbO  is  formed  by  dissolv- 
ing lead  or  its  oxide  in  HN03. 

Lead  Chromate  PbCr04— 

A yellow  precipitate,  formed  when  a soluble  chromate  reacts 
on  a soluble  Pb  salt — it  is  the  test  for  either — it  is  used  as  a pig- 
ment and  is  highly  poisonous  from  both  of  its  ingredients. 

Sulphide  PbS — 

“Galena”  is  the  chief  ore  of  lead — occurs  in  cubical  crystals 
with  metallic  lustre — is  precipitated  by  H2S  from  Pb  salt. 


White  Lead — 

Is  one  of  the  oldest  pigments— A basic  lead  carbonate  of 
variable  composition,  made  by  passing  CO2  through  solution 
of  Pb(C2H302)2  (French  method),  or  exposing  rolls  of  Pb  to 
atmosphere  of  H(C2H302)  and  C02— often  adulterated  with 
BaSC>4  and  CaSO*. 

LEAD  SALTS. 


NaOH : precipitates  Pb20(0H)2,  soluble  in  excess. 

H&NOH: “ Pb20(0H)2,  insoluble  in  excess. 

Na2C02: “ white  PbC03,  soluble  in  excess. 

H2S : “ black  PbS. 

KI: “ yellow  PbL,  soluble  in  hot  H2O. 

K2CrO 4: “ yellow  PbCr04. 

HC1: “ PbCb,  in  strong  solutions. 

H2SO±: “ PbS04,  insoluble  in  H2O  or  dilute  H2SCU— soluble 

in  H2S04  concentrated. 


H^NC2H202\  decomposes  PbS04. 

Fe  and  Zn  : ...  separate  crystalline  Pb  from  solutions. 

Blowpipe : ....  gives  a yellow  coating  and  malleable  lead. 

THALLIUM— Te. 

At.  Wt.  204.  Val.  I or  III. 

A spectrum  metal  found  by  Crookes  in  the  mud  from  Swed- 
ish H2SO4  chambers.  Occurs  in  minute  quantities  in  combina- 
tion with  other  metals — chemical  character  is  partly  like  Na, 
partly  like  Pb  Oxidizes  rapidly  in  the  air,  forming  Te20  and 
and  Te203 — most  salts  are  soluble. 

ALUMINIUM — Al. 

At.  Wt.  274.  Yal.  III.  or  IV. 

History — 

In  1817  the  metal  was  first  obtained  by  Wohler  from  Al3 
Cl6  and  K. 

Occurrence — 

Widely  distributed  in  large  quantities  as  the  oxide,  in  ruby, 
sapphire,  corundum  and  emery — as  silicate,  in  clay,  mica  and 
most  crytalline  rocks. 

Preparation — 

(1)  By  fusing  the  chloride  with  metallic  Na. 

(2)  Electrolysis  of  the  chloride,  AI2CI6. 


83 


Properties — 

Silver  white,  highly  malleable,  ductile  and  sonorous — a 
light,  tenacious  metal— in  bulk,  A1  is  stable  in  ,the  air — easily 
soluble  in  HC1,  or  boiling  H2SO4,  not  in  HNO3.  NaOH  dissolves 
it  forming  an  aluminiate. 

A1  + 3NaOH  = Na3A103  + 3H. 

Unites  directly  with  S group,  P,  As,  Si,  and  Haloids.  In  foil, 
A1  burns  in  the  air,  and  decomposes  H2O.  At  100°  A1  alloys 
easily  with  some  metals,  especially  with  Cu  in  aluminium 
bronze. 

Metallic  A1  is  used  in  scientific  instruments,  watch-springs, 
ornaments,  etc. — Aluminium  bronze,  in  watches,  spoons,  etc. — 
Aluminium  silicate  or  clay,  in  pottery,  bricks,  etc. — Aluminium 
acetate,  as  a mordant  in  dyeing  cotton  cloth. 

Al  COMPOUNDS. 

Aluminium  Oxide  AI2O3  “Alumina” — 

Found  in  nature  as  ruby,  sapphire,  etc. — prepared  by  heating 
the  hydrate  Al2(OH)e  or  igniting  an  alum — The  oxide  is  infus- 
ible, save  in  oxy hydrogen  flame — if  ignited,  is  insoluble  in  H2O 
and  acids — before  ignition,  is  soluble  in  acids,  and  KOH. 

Aluminium  Hydrate  A12(OH)g — 

A voluminous,  white  precipitate,  formed  when  (H4N)0H  is 
added  to  solution  of  an  Al  salt — If  any  vegetable  coloring  mat- 
ter should  be  present,  it  would  be  precipitated  with  the  hydrate 
forming  an  insolublecompound,  technically  known  as  a “Lake” 
— Al2(OH)e  freshly  formed,  dissolves  in  excess  of  KOH  or  in 
acids.  After  long  standing  its  solubility  diminishes — heated,  it 
forms  AI2O3. 

Aluminium  Chloride  ALCle — 

Formed  when  Cl  gas  is  passed  over  hot  Al — the  hydrated 
salt  is  formed  by  dissolving  Al2(OH)6  in  HC1 — ALCle  is  a color- 
less deliquescent  salt — forms  double  salts  with  other  metals — 
aqueous  solution  is  a disinfectant. 

Aluminium  Sulphate  Al2(S04)3  * I8H2O — 

A white  crystalline  solid,  obtained  from  solution  of  Al2(OH)6 
in  H2SO4,  or  industrially,  from  clay  and  H2SO4 — easily  soluble 


84 


in  water— becomes  anhydrous  with  heat — at  high  temperature 
decomposes  to  AI2O3  + SO4. 

Alums — 

Aluminium  sulphate  combines  with  the  alkali  sulphates 
to  form  “alums.”  The  chief  alum  is  potassium  aluminium 
sulphate  K2  AI3  (SCLH  + 24Hl> 0,  “common  alum.”  Alum 
is  a general  term,  applied  to  a series  of  double  salts  of  similar 
composition  and  crystalline  form.  Iron,  chromium  and  man- 
ganese form  similar  derivatives  by  replacing  the  AI2 — a further 
series  is  formed,  by  replacing  K2  with  Na2  or  (H4N)2,  thus  we 
have  iron  alum.  Fe2K2(S04)4  + 24H2O,  chrome  alum  Cr2K2- 
(S04)4  + 24H2O,  and  others  of  similar  formula. 

ALUMINIUM  SALTS. 

The  haloids,  sulphates,  nitrates  and  acetates  are  soluble. 

NaOH  or  KOH : gives  Al2(0H)6  soluble  in  excess. 

H±NOH\  gives  Al2(0H)6  insoluble  in  excess. 

NazCOs : “ “ 

(H4W)2S  “ “ “ “ “ 

Al  silicates. 

Occur  ranee — 

Granite  contains  both  Al  silicate  and  alkaline  silicates.  When 
it  disintegrates,  the  alkaline  silicates  are  washed  away,  leaving 
insoluble  aluminum  silicate,  clay — purest  form  of  clay  in  Kao- 
lin— Topaz,  Beryl  and  Lapis  Lazuli,  are  important  silicates,  es- 
pecially the  latter,  of  which  the  powder  was  the  old  source  of 
Ultramarine. 

Ultramarine — 

Is  now  made  by  heating  a mixture  of  clay,  Na2C03,  S and 
C. — this  gives  a colorless  compound,  turning  green — by  gently 
heating  with  sulphur,  it  forms  the  blue  variety — variation  in 
heat  and  ingredients  gives  a variet\r  of  colors. 

Porcelain — 

Is  essentially  an  Al  glass  made  from  kaolin — true  porcelain 
is  translucent — slightly  attacked  by  reagents. 

Pottery— 

Is  an  impure  porcelain  made  from  clay — its  red  color  due 
to  FeO. 


— 85- 


GLASS. 

Properties — 

A mixture  of  silicates  of  the  alkalies  and  other  metals — a 
thin  liquid  at  high  temperatures,  becoming  viscous  as  it  cools. 
Amorphous,  transparent  solid,  little  affected  by  acids  or  water. 
No  single  silicate  has  all  these  properties.  The  alkali  silicates 
are  amorphous  and  transparent  but  soluble.  Others  are  in- 
soluble, but  crystalline.  A combination  has  the  required  prop- 
erties. There  are  four  classes  of  glass. 

I —  Bohemian  or  Hard  Glass — 

K and  Ca  silicate,  which  melts  at  high  temperature  and  re- 
sists reagents — Used  in  laboratories. 

II—  French  Glass— “Crown  Glass”— 

Na  and  Ca  silicate  of  blue-green  color,  harder  than  Bohe- 
mian but  easier  melted.  Another  name  is  “crown  glass.” 

III—  Bottle  Glass — 

Impure  glass,  colored  by  Fe203 — a silicate  of  the  alkalies, 
Ca,  Mg  and  Al. 

IV —  Lead  or  Crystal  Glass— 

K and  Pb  silicates — Pb  increases  refraction,  hence  used  as 
“strass”  glass  to  imitate  jewels.  Used  in  optics  as  “flint” 
glass. 

Colored  Glass — 

Contains  Cu20(red),  or  Au(ruby),  CuO(green),  Co(blue), 
Mn( violet),  etc. 

GALLIUM  Ga.  INDIUM  In. 

At.  Wt.  70.  Val.  Ill  or  IV.  At.  Wt.  113.4.  Val.  Ill  or  IV. 

Two  metals  discovered  through  the  spectroscope,  and  close- 
ly analogous  to  aluminium — the  aluminium  group,  Al,  Ga,  In, 
are  true  metals — their  oxide  M2O3,  is  a weak  base,  acting  as  an 
acid  in  presence  of  strong  alkalies — the  metals  do  not  decom- 
pose water — are  easily  dissolved  by  haloid  acids  with  evolution 
of  H.  -Their  valence  is  III.  or  IV. 

AI2O3  Ga2C>3  I112O3 

AI2CI6  Ga2Cl6  Iti2  Cle 

Alums  are  easily  formed  from  their  sulphates. 


-86 


MANGANESE-Mn. 

At.  Wt.  55.  Val.  II.  and  IV. 

Occurrence — 

Never  free — chief  one  is  Pyrolusite,  (M11O2) 

Extraction — 

By  igniting  the  oxide  with  C. 

Properties— 

Hard,  gray,  metal — fuses  with  difficulty — not  stable  in  the 
air — metal  has  no  technical  use — forms  an  alloy  with  Fe  which 
is  used  in  Bessemer  process. 

OXIDES  OF  MANGANESE. 

MnO,  Mn203,  Mn304,  Mn02,  Mn207. 

Manganous  Oxide  MnO — 

Formed  by  reduction  of  higher  oxides  or  by  ignition  of  the 
oxalate — reduced  by  C to  the  metal — MnO  is  the  basis  of  com- 
mon manganese  salts— its  hydrate  Mn(OH)2  is  a white  precip- 
itate easily  oxidized— like  Mg(OH)2  it  is  not  precipitated  in 
H4NCI  solution. 

Manganic  Oxide  Mn203~ 

A black  powder — made  by  igniting  the  other  oxides  in  oxy- 
gen— its  hydroxide,  Mn(OH)3  is  prepared  from  Mn(OH)2. 
Treated  with  dilute  acids,  the  oxide  or  hydroxide  gives  mangan- 
ous salts. 

Mn203  + H2SO4  — MnS04  + Mn02  + H20. 

M/32O3  occurs  as  mineral  “braunite.” 

Mangano=Hanganic  Oxide  Mn304  or  MnO  * Mn2Os— 

Found  as  a red-brown  powder,  by  ignition  of  the  carbon- 
ate— acts  like  a mixture  of  MnO  + M^Os — is  isomorphous 
with  magnetite,  Fe304. 

M3O4  occurs  as  mineral  “hausmanite.” 
rianganese  Dioxide  Mn02 — 

Occurs  as  “pyrolusite’ ’ — made  by  igniting  manganous 
nitrate  Mn(N03)2~ Mn02  is  a black  powder,  which,  heated, 
yields  O — treated  with  HC1,  gives  off  chlorine — its  hydroxide  is 
Mn(OH)2,  from  which  the  unstable  manganites  are  derived. 


OTHER  Mn  COMPOUNDS. 


Hanganous  Ammonium  Phosphate  M11H4NPO4 — 

Like  the  Mg  salt  is  formed  when  H4NOH  and  alkaline  phos- 
phate are  added  to  a Mn  salt. 

Manganous  Carbonate  MnCOs — 

Precipitated  by  alkaline  carbonates  from  manganous  salts — 
a white  powder  which  easily  oxidizes  turning  brown — like 
CaC03,  is  soluble  in  natural  waters  which  contain  CO2. 

Manganous  Sulphide  MnS — 

Precipitated  from  manganous  solution  by  the  alkaline  sul- 
phides— a flesh  colored  precipitate  turning  brown,  then  green. 
MnS2  corresponding  to  Mn02  is  likewise  found  in  nature  as 
“hauerite.” 

Mn  ACIDS. 

Manganic  Acid  H2Mn04— 

The  oxide  Mn(>3  is  not  known — The  hydrated  oxide,  man- 
ganic acid  H2MnC>4  breaks  up  in  solution  to  permanganic  acid 
and  Mn02. 

3H2Mn04  = 2HM11O4  + M«02  + 2H20 
Manganic  acid  is  dibasic,  corresponding  to  H2SO4  or  H2C2O4 
and  the  salts  of  these  three  acids  are  isomorphous. 

Potassium  Manganate  K^MnCL — 

The  K salt  is  formed  from  Mn02+K0H. 

3Mn02  + 2K0H  = K2Mn04  + Mn203  + H20. 

Manganates  are  of  green  color  and  easily  decomposed  by 
acids  to  permanganates. 

Permanganic  Acid  HMnCL — 

Known  both  in  salts  and  free  state — the  free  acid  is  made 
from  its  Ba  salt.  HMnC>4  is  a powerful  oxidizing  agent — de- 
composes above  40° — salts  are  purple  and  generally  made  by 
decomposition  of  manganates  with  acids. 

Potassium  Permanganate  KMnCL  (K2Mn208)  — 

Chief  salt  of  permanganic  acid,  is  formed  when  CO2  acts  on 
K manganate,  till  the  green  color  changes  to  red. 

3K2Mn04  + 2C02  = 2KMn04  + 2K2C03  + M«02. 

KMnC>4  crystallizes  in  dark  red  prisms — isomorphous  with 
KCIO4. 


88 


A powerful  oxidizing  agent — in  oxidizing  it  is  reduced  to  the 
colorless  manganate,  hence  the  persistence  of  the  permanganate 
color  tells  when  an  oxidation  is  complete — Na-permanganate 
is  used  commercially  as  a disinfectant  under  name  of  Condy’s 
fluid. 

OTHER  Mn  COMPOUNDS. 

flanganous  Chloride  MnCL — 

Red  deliquescent  crystals — formed  when  Cl  is  made  from 
Mn02  and  HC1. 

Mn02  + 4HC1  = MnCl2  + Cl2  + 2H20. 

Manganous  Sulphate  MnS04 — 

Bright  red  or  pink  crystals,  formed  from  H2SO4  and  MnO. 

MnO  + H2S04  = M11SO4  + H20. 

If  higher  oxides  are  used,  oxygen  will  separate. 

MANGANOUS  SALTS. 

Color — reddish  when  crystalline — otherwise  colorless. 

NaOH  and  KOH  give  white  Mn(OH)2  becoming  brown. 


iYa2COs gives  “ “ “ “ 

BaCOs precipitates  “ only  from  MnS04. 

(H4AT)2S “ MnS,  flesh-colored,  turning  green. 

H2S “ MnS  but  slightly,  in  neutral  solution. 

KNOs  and 

iVa2C03 heated  with  -ous  salts  give  K2Mn04  (characteristic). 

Na<2,B±Oi “ “ “ amvthyst  color. 


IRON— Fe. 

At.  Wt.  56.  Val  II.  and  III. 

History — 

The  use  of  iron  dates  back  to  the  early  legends,  when  the 
source  was  probably  meteoric.  Later  it  was  obtained  from 
ores,  as  the  great  slag  fields  of  India  give  witness. 

Occurrence — 

The  most  widely  distributed  metal — seldom  found  native 
save  in  meteors,  where  it  generally  contains  some  nickel — occurs 
in  the  blood  and  mineral  waters — chief  ores  are  magnetite  Fe3 
O4,  hematite  Fe203,  limonite  Fe203  +H20,and  siderite  FeC03. 


89 


IRON  ORES. 

Hagnetite  Fes04—  (Fe203,  FeO)— 

Is  tlie  richest  and  best  of  iron  ores — occurs  in  crystalline  and 
amorphous  forms — is  reduced  with  difficulty,  but  yields  excellent 
iron — magnetite  may  have  Fe20s  replaced  by  Mn203,  and  FeO 
by  ZnO — it  is  then  an  ore  of  zinc  called  franklinite — or  may 
have  its  Fe203  replaced  by  Cr203,  and  is  then  chromite. 

Hematite  Fe203 — 

One  of  the  most  common  ores  of  iron — occurs  also  in  a crys- 
talline modification  known  as  specular  iron  ore. 

Limonite — 

Or  bog  iron  is  a mixture  of  Fe203  and  Fe2(OH)6 — yellow 
ochre  is  the  clayey  variety. 

Siderite — 

Known  also  as  spathic  iron  ore,  and  when  mixed  with  clay 
as  clay  iron  stone,  is  the  chief  English  ore. 

Extraction — 

Ores  are  pulverized  and  roasted,  to  drive  out  water,  car- 
bonic acid,  or  sulphur,  and  to  convert  oxides  to  ferric  oxide 
Fe20s,  which  easily  reduces.  Ores  are  then  heated  with  C and 
fluxes  in  the  blast  furnace.  The  furnace  is  charged  with  alter- 
nate layers  of  fuel  (charcoal,  coke  or  anthracite),  ore  and  flux — 
a powerful  blast  of  hot  air  oxidizes  the  C to  CO2 — this  is  re- 
duced by  hot  C farther  up  the  pile  to  CO — CO  passing  through 
the  hot  ore  reduces  Fe203  to  a spongy  mass  of  metallic  Fe, 
which,  mixed  with  flux  and  earthy  impurities,  settles  to  hotter 
part  of  furnace.  Here  the  iron  forms  a fusible  compound  with 
C,  and  drops  to  bottom  of  furnace  to  be  drawn  off— the  flux  and 
impurities  melt  to  a liquid  “slag”  which  floats  above  the  molten 
iron. 

VARIETIES  OF  IRON. 

Pig=Iron  or  Cast=Iron 

is  the  crude  iron  as  obtained  from  the  furnace  and  may  contain 
carbon,  phosphorus,  sulphur,  silicon,  manganese,  etc.  If  cooled 
rapidly  when  taken  from  the  furnace  most  of  its  carbon  remains 
in  combination  and  it  is  then  known  as  white  cast  iron — this 
leaves  no  residue  when  dissolved  in  acid,  as  all  the  C unites 


90 


with  the  H set  free.  If  theiron  is  cooled  slowly  theC  separates 
as  graphite  and  is  so  left  when  iron  is  dissolved — this  variety  is 
gray  cast  iron.  If  the  ore  contains  much  manganese,  this  is 
renuced  at  same  timeand  such  manganese-containing  iron,  may 
combine  with  a greater  auantity  of  carbon,  and  is  known  as 
spiegel  iron.  The  presence  of  C,  Si  and  P renders  cast  iron 
brittle,  hence  they  must  be  removed  before  it  can  be  welded — 
these  substances  are  oxidized  in  the  puddling  furnance,  which 
yields  wrought  iron  containing  less  than  .6  per  cent  of  C. 

Steel — 

Is  intermediate  between  cast  and  wrought  iron — may  be 
made  by  adding  C directly  to  wrought  iron  as  in  the  Cementa- 
tion process,  which  produces  very  good  steel,  or  in  Bessemer 
process,  by  decarbonizing  cast  iron  to  form  wrought  iron,  then 
adding  enough  cast  iron  to  produce  the  desired  percentage  of 
C for  steel. 

Properties — 

Commercial  iron  is  never  pure. — The  several  varieties  differ 
chiefly  by  reason  of  impurities,  especially  C as  given  above — 
purest  form  is  wrought  iron,  especially  in  piano  wire  which  con- 
tains but  three  per  cent,  of  impurities. 

Pure  Fe  may  be  obtained  by  igniting  the  oxide  in  a current 
of  H — it  is  then  a silver  white  metal — softer  and  more  malleable 
than  wrought  iron,  but  less  tenacious. 

Iron  oxidizes  easily  in  moist  air  or  under  water  to  Fe3C>4 — 
finely  divided  it  decomposes  H2O  at  100° — dilute  acids  dissolve 
Fe  with  evolution  of  H — concentrated  HNO3  does  not  affect  it, 
due  probably  to  thin  coat  of  the  oxide — CO2  in  water  forms 
FeCOs — Fe  unites  directly  with  haloids  with  S,  C,  Si,  and  the 
metals. 

OXIDES  OF  Fe. 

Iron  forms  four  oxides  FeO , Fe^O^,  FesO±,  and  the  hypo- 
thetical FeOz.  FeO  and  Fe203  are  basic  and  form  salts — FeO,  a 
strong  base,  corresponds  to  CaO — Fe203,  a weak  base,  corre- 
sponds to  AI2O3 — Fe03,  like  Mn03,  exists  only  in  the  K salt. 

Ferrous  Oxide  FeO — 

Does  not  occur  free  but  always  as  a salt,  in  nature — difficult 
to  make,  as  it  absorbs  0 — may  be  formed  by  reducing  Fe203 
with  H — a black  powder. 


91 


Ferrous  Hydrate  Fe(OH)2 — 

The  base  of  ferrous  (Fe)  salts — a white  precipitate  thrown 
out  of  ferrous  salts  by  action  of  the  alkalies  in  a current  of  H. 
On  exposure  to  air  it  oxidizes  to  green,  then  red-brown  color — 
insoluble  in  water^-a  strong  base. 

FERROUS  SALTS— Fe. 

White  when  anhydrous — blue-green  when  hydrated. 

KOH  and  NaOH  form  white  Fe(0H)2,  easily  oxidized. 


K2CO3 precipitates  white  FeCOs,  easily  oxidized. 

H2S “ black  FeS,  only  in  dilute  neutral  solutions. 

(H4AO2S “ “ FeS,  soluble  in  acids. 

K^FcCqNq “ white  Fe^FeCeNe,  turns  blue  by  oxidation. 

KqFc2C\2N\2 ' “ blue  Fe3Fe2Ci2Afi2  “Turnbulls  blue.” 

KCNS gives  no  color  when  no  Fe2  salts  are  present. 

Na2FLFO± precipitates  white  Fe3(P04)2,  turns  blue  by  oxidation. 

Tannic  acid gives  no  color  when  no  Fe2  salts  are  present. 

Na2B^07 gives  a colorless  or  yellow  bead. 


Ferrous  salts  absorb  NO  in  the  cold,  turning  black,  and  are 
used  as  a reagent  for  HNO3. 

Ferric  Oxide  Fe203 — 

Occurs  in  nature  as  hematite — prepared  by  ignition  of  fer- 
rous sulphate — 

2 FeS04  = Fe203  + S03  + S02. 

Thus  prepared  Fe203  is  a dark  red  powder,  used  as  a paint 
under  name  “Colcothar” — Fe203  is  a very  hydroscopic  sub- 
stance and  like  Mn02  aids  in  preparation  of  0 from  KCIO3. 

Ferric  Hydrate  Fe2(OH)6— 

The  basis  of  ferric  salts — red  brown  voluminous  precipitate, 
made  by  treating  a ferric  salt  with  an  alkali  hydrate — red  heat 
renders  it  insoluble  in  acids.. 

FERRIC  SALTS. 

Color — 

If  anhydrous  the  neutral  (Fe2)  salts  are  without  color — 
when  crystallized  they  are  yellow  or  reddish  yellow,  save  the 
nitrate  and  fluoride  which  are  colorless,  and  the  acetate  and 
thiocyanate  which  are  deep  red — All  ferric  salts  are  reduced  to 
ferrous  by  H2S. 


KOII,  K2COs\  precipitates  red  brown  Fe2(OH)6  soluble  in  acids— this  pre- 
and  BnCO$:  j cipitation  is  prevented  by  sugar. 


HzS: reduces  ferric  to  ferrous  salts,  setting  free  S. 

Fe2Cl6  + H2S  — 2 FeCl2  + 2 HC1  + S. 

(H4N)2S: gives  FeS. 

Fe2Cl6  + (H4N)2S  = 2 FeS  + 6 H4NC1  + S. 

K^FeCeNe  : gives  blue  Fe2FeCeN6,  “Prussian  blue.” 

KCNS: gives  red  Fe2(CNS)6,  organic  acids  hinder. 

Tannic  acid:....  gives  black  iron  tannate  (ink). 


Ferric  Acid  FUFeCU — 

Corresponds  to  H2SO4  and  H2Mn04 — free  acid  is  not  known 
as  it  is  unstable.  Potassium  ferrate  K2Fe04  is  formed 
when  iron  filings  are  fused  with  KNO3 — (see  K^MnCU) — crys- 
tals of  K salt  form  in  dark  red  prisms. 

Ferrous  Chloride  FeCU — 

The  white  anhydrous  salt  is  formed  when  HC1  gas  passes 
over  excess  of  heated  iron— green  deliquescent  FeC^  is  formed 
by  the  solution  of  Fe  in  HC1 — oxidizes  easily  to  Fe2Cl6 — all 
ferrous  salts  absorb  NO. 

Ferrous  Bromide  FeBr2 — 

Made  by  action  of  Br  on  excess  of  Fe  in  the  presence  of 
H2O — used  in  preparation  of  KBr. 

Ferrous  Sulphate  FeS04*7H20 — 

“Green  vitriol”  is  a bi-product  of  Cu  works — obtained  pure 
from  FeS  and  H2SO4. 

FeS  + H2S04  = FeS04  + H2S. 

Crystals  effloresce  in  dry  air,  and  if  heated  lose  six  molecules 
of  water  to  form  the  colorless  salt  FeSC^'H^O — at  still  greater 
heat  FeS04  decomposes. 

2FeS04  = Fe203  + S03  + S02. 

From  this  reaction  the  “fuming”  or  “Nordhausen”  sulphuric 
acid  is  obtained. 

Ammonium  Ferrous  Sulphate  FeSC>4‘  (H4N)2S04.6H20 — 

A stable  salt,  prepared  bv  evaporating  together  equal  parts 
of  FeS04-7H20  and  (H4N)2S04. 

Ferrous  Nitrate  Fe(N0s)2 — 

Unstable,  green  salt,  made  by  action  of  HNO3  on  excess  of 
iron,  or  from  Ba(N03)2  and  FeSC>4. 

Ba(N03)2  + FeS04  = Fe(N03)2  + BaS04. 


93 


Ferric  Chloride  Fe2Clc> — 

Anhydrous,  by  heating  iron  in  current  of  Cl — hydrated  Fe2- 
Cle  ' 6H2O  bypassing  Cl  gas  into  solution  of  FeCb—anhydrous 
is  brown — hydrated,  red — both  are  deliquescent — soluble — vola- 
tile at  high  heat — forms  basic  salts  easily — used  as  a reagent  for 
organic  substances. 

Other  (Fe)2  salts  are  formed  by  oxidation  of  the  (Fe)  fer- 
rous. 

CYANOGEN  COMPOUNDS  OF  IRON. 

Ferrous  Cyanide  Fe(CN)2 — 

Is  found  as  a white  unstable  precipitate  when  KCN  is  added 
to  the  solution  of  a ferrous  salt — with  an  excess  of  KCN  the 
double  salt  Fe(CN)2,  4KCN  or  K^FeCeNe  is  formed. 

Potassium  Ferrocyanide  ^FeCeNe — 

This  is  known  commercially  as  yellow  prussiate  of  potash, 
and  made  by  heating  K2CO3  + Fe  with  nitrogenous  bodies  to  a 
red  heat — a sweetish,  non-poisonous  salt — used  in  dyeing  and 
as  a test  for  Fe2  and  CN  salts — with  concentrated  HC1  it  forms 
the  acid,  hydrogen  ferrocyanide,  or  ferrohydrocyanic  acid  H4 
Fe(CN)6,  which  is  unstable — with  ferric  salts,  ferrocyanide 
forms  prussian  blue. 

Potassium  Ferricyanide  KeFe2Ci2Ni2 — 

Is  formed  when  the  ferrocyanide  is  treated  with  Cl. 

2K4Fe(CN)6  + Cl3  = K6Fe2(CN)i2  + 2KC1. 

Separates  in  red  prisms,  known  as  red  prussiate  of  potash — 
its  water,  solution  reduces  to  ferrocyanide  by  light — its  alkali 
solution  is  a strong  oxidizing  agent — with  ferrous  salts,  ferricy- 
anide forms  “Turnbulls  blue.” 

IRON  SULPHIDES. 

Ferrous  Sulphide  FeS — 

Gray,  metallic  mass,  made  by  fusing  Fe  and  S — or  merely 
moistening  a mixture  of  Fe  and  S at  ordinary  temperature. — Is 
obtained  pure  through  precipitation  from  ferrous  salts  by 
(H4N)2S — FeS  is  soluble  in  acid  and  largely  used  for  making 
H2S. 

FeS  + 2HC1  = H2S  + FeCl2. 


94 

Ferric  Sulphide  Fe2S3 — 

Black  sulphide  analogous  to  Fe2C>3,  made  by  heating  Fe 
and  S in  proper  proportions. 

“Pyrite”  FeS2— 

A natural  sulphide  not  analogous  to  any  known  oxide. 
NICKEL— Ni. 

At.  Wt.  58.6— Val.  II. 

Occurrence — 

Generally  accompanies  Co — chief  ore  is  NiAs“Kupfernickel” 
— with  Fe  only  in  meteorites. 

Properties — 

Hard  white  metal — malleable— weldable — ductile — more  te- 
nacious than  iron — magnetic — but  slightly  affected  by  the  air, 
hence  used  to  plate  other  metals — soluble  in  mcst  acids  but  like 
Fe,  is  passive  in  HNO3  concentrated. 

Oxides  and  Salts — 

Of  nickel  are  closely  analogous  to  those  of  Co. — Salts  are 
green,  complementary  to  the  red  of  Co  salts — the  hydrate  Ni02 
H2  is  somewhat  soluble  in  H4NOH,  with  a blue  color  similar  to 
Cu. 


COBALT-Co. 

At.  Wt.  58.6.  Val.  II  or  III. 

History — 

The  name  “kobold”  was  given  first  to  minerals  which  col- 
ored glass  blue,  afterwards  to  the  metal  which  was  discov- 
ered i i them. 

Occurrence — 

Never  in  the  free  state,  usually  in  combination  with  As  and 
S— chief  ores  are  “smaltite,”  C0AS2,  and  “cobaltite,”  CoAs2* 
C0S2. 

Extraction — 

By  roasting  the  ore  a crude  arseniate  is  obtained,  “zaffre 
this  treated  with  KHSO4  gives  the  crude  oxide  which  is  reduced 
with  carbon  to  metallic  Co. 


Properties — 

A reddish  metal — lustrous — malleable— ductile — difficultly 
fusible — harder  and  more  tenacious  than  iron — can  be  magnet- 
ized—not  attacked  by  air  or  water — slightly  by  HClandH2S04 
— easily  by  HNO3.  Chief  compounds  are  cobaltous,  i.  e.,  where 
Co  is  bivalent. 


COBALT  OXIDES. 

CobaltQus  Oxide  CoO — 

Cobalt  forms  three  oxides,  CoO,  C02O3,  C03O4.  CoO  is  a 
brown  powder  formed  by  heating  C02O3  with  H — is  the  basis 
of  cobaltus  salts — with  water  it  forms  Co(OH) 2,  Cobaltous 
hydrate,  a rose  colored  precipitate  turning  brown — Co  (OH)  2 
like  Fe02H2  must  be  made  in  an  atmosphere  of  H. 

Cobaltic  Oxide  C02O3 — 

A black  powder  formed  by  gentle  ignition  of  Co(N03)2 — it 
is  basis  of  cobaltic  salts,  which  are  ver}r  unstable. 

Cobaltous=Cobaltic  Oxide  C03O4 — 

A black  pow  der  formed  by  strong  ignition  of  the  nitrate — 
it  is  not  magnetic  and  forms  no  salts. 

OTHER  Co  COMPOUNDS. 

Cobaltous  Chloride  C0CI2 — 

Obtained  by  dissolving  the  oxide  or  carbonate  in  HC1 — it 
crystallizes  with  6H2O  to  form  red  crystals,  which,  when  anhy- 
drous are  blue — the  red  crystals  are  also  turned  blue  by  H2SO4, 
which  takes  up  its  water  of  crystallization — characters  written 
with  the  red  solution  are  almost  invisible  but  turn  to  distinct 
blue  when  heated  (sympathetic  ink).  The  Br  and  I compounds 
are  similar. 

Cobaltous  Sulphate  C0SO4  7H2O — 

Crystallizes  in  dark  red  prisms,  and  resembles  ferrous 
sulphate. 

Cobaltous  Nitrate  Co(NOs)2 ' 6 H20 — 

Forms  in  red  deliquescent  prisms — is  much  used  in  blowpipe 
work,  and  gives  characteristic  blowpipe  reaction swith  Zn  (Rin- 


maus  green)  and  A1  (Thenard’s  blue)  also  a blue  with  phos- 
phates. 

Cobalt  Silicate — 

No  silicates  are  found  in  nature,  but  when  glass  is  fused  with 
a Co  salt  it  forms  a dark  blue  silicate,  called  “smalt” — when 
reduced  to  powder  is  used  as  a paint. 

COBALT  CYANIDES. 

Cobalt  Cyanides — 

Are  formed  by  dissolving  the  hydrate  in  KCn — in  excess  of 
KCN  a double  salt  is  formed,  Co(CN)2.2KCn — when  this  is 
boiled  with  an  oxidizing  agent  a cobalticyanide,  KeCo2(CN  )i2, 
is  formed  and  from  this  KOH  does  not  precipitate  Co2(0H)e. 

COBALT  SULPHIDES. 

Cobalt  Sulphides  CoS— 

The  most  common  is  CoS,  a black  precipitate,  formed  when 
(H4N)2Sis  added  to  a Co  salt — insoluble  in  alkalies  or  cold 
dilute  HC1 — soluble  in  strong  acids. 

GOLD— Au. 

At.  Wt.  196-7  Val.  1.  and  III. 

History — 

Long  known  to  the  ancients  and  called  the  king  of  metals — 
chief  object  of  alchemists,  was  the  transmutation  of  base  metals 
to  gold. 

Occurrence — 

Chiefly  as  native  gold,  alloyed  with  more  or  less  Ag,Cu,Pb 
and  Bi — is  widely  distributed,  but  generally  occurs  in  small 
quantities — its  ores  are  unimportant. 

Properties — 

Most  malleable  and  ductile  of  metals — highlj^  tenacious — at 
high  temperature  is  volatile — melted  has  a red  yellow  color — 
on  cooling,  Au  contracts  more  than  other  metals — no  single  acid 
dissolves  Au,  but  with  aqua  regia  (HC1  + HNO3)  it  forms  the 
chloride — the  other  haloids  also  attack  it,  and  melted  KNO3 
forms  the  oxide. 


97 


Auric  Chloride  A11CI3 — 

Formed  by  solution  of  Au  in  Aqua  regia — a reddish  brown, 
deliquescent  mass  of  crystals,  which  dissolve  readily  in  alcohol 
and  ether — All  reducing  agents  are  oxidized  by  AuCU,  with 
precipitates  of  gold — thus  FeSC>4  precipitates  gold  lrom  the 
chloride  as  a lustreless  brown  precipitate — and  stannous  chlor- 
ide SnCU,  precipitates  the  oxide  AU2O  as  “purple  of  Cassius.” 

Aurous  Chloride  AuCl. 

When  auric  chloride  is  heated  to  180°  it  forms  aurous 
chloride,  AuCl,  a white  powder  in  H2O— ignited  AuCl  decom- 
poses to  Au  + CL 

Aurous  Oxide  AU2O — 

Formed  in  “purple  of  cassius”  also  when  KOH  acts  on 
aurous  chloride— dark  violet  powder — decomposed  by  heat — 
changed  to  Xu  and  AUCI3  by  HC1. 

Auric  Oxide  AU2O3 — 

When  A-UCI3  is  heated  with  MgO  a brown  precipitate  is 
formed.  If  excess  of  MgO  in  is  removed  by  concentrated  HNO3 
auric  oxide  is  left  as  a brown  powder. 

Auric  Hydroxide — 

If  MgO  is  removed  with  dilute  HNO3,  Au(OH)3  remains  as 
a yellow-red  powder — oxide  and  hydroxide  are  insoluble  in 
water  and  acids,  but  have  acid  properties  and  are  soluble  in 
alkalies.  The  hydroxide  is  called  auric  acid  and  forms  aurates, 
which  are  derived  from  meta-auric  acid,  HAUO2. 

Auric  Sulphide  AU2S3 — 

H2S  or  soluble  sulphides  precipitate  auric  sulphide,  as  a 
brown  precipitate,  soluble  in  alkaline  sulphides. 

Gold  Cyanide  AuCN — 

When  gold  or  its  oxide  is  dissolved  in  KCN  the  colorless 
double  salt,  AuCN . KCN,  may  be  crystalized  out — from  this 
metallic  Au  is  easily  precipitated  by  electrolysis,  hence  used  in 
gilding. 


98 


PLATINUM— Pt. 

At.  Wt.  194.4.  Val.  II.  and  IV. 

Occurrence — 

Found  only  in  metallic  state,  generally  alloyed  with  pal- 
ladium, iridium  and  other  rare  metals. 

Extraction — 

The  ore  is  dissolved  in  aqua  regia,  and  precipitated  by  (H4 
N)C1 — this  forms  by  ignition  a double  salt  of  iridium  bearing 
platinum  in  a spongy  mass  ( platinum  sponge.)  This  is  used 
directly  in  making  platinum  vessels— if  desired  the  iridium, 
which  diminishes  malleability  but  increases  hardness  and  re- 
sistance to  reagents,  may  be  farther  removed. 

Properties — 

Heavy,  lustrous,  gray-white  metal — softer  than  Ag — not 
affected  by  air  nor  common  acids — dissolved  by  aqua  regia  and 
chlorine  water — attacked  by  fusing  caustic  alkalies — finely 
divided  as  in  Pt  sponge,  and  Pt  black,  it  remarkably  influences 
the  chemical  combination  of  gases. 

PLATIUM  COMPOUNDS. 

Platinic  Chloride  PtCL — 

Is  the  most  important  salt,  made  by  dissolving  Pt  in  aqua- 
regia.  It  dissolves  in  water  to  a red  yellow  solution,  important 
because  it  forms  with  H±N  or  K an  insoluble  salt — the  oxides 
and  salts  of  Pt  in  general  are  formed,  and  act  in  the  same  way 
as  those  of  Au. — Their  formula  is  different  because  Pt  acts  with 
valence  of  two  or  four,  thus:  PtCl2,  PtCU,  PtS2,etc — all  Pt  salts 
leave,  on  ignition,  a residue  of  metallic  Pt. 


